EPA-600/3-76-108
November 1976
ATMOSPHERIC FREONS AND HALOGENATED COMPOUNDS
BY
Alan Appleby
Rutgers University
Department of Environmental Science
New Brunswick, New Jersey
08903
R-800833
Project Officers
Joseph J. Bufalini and Bruce W. Gay, Jr.
Gas Kinetics and Photochemistry Branch
Environmental Sciences Research Laboratory
Research Triangle Park, North Carolina 27711
ENVIRONMENTAL SCIENCES RESEARCH LABORATORY
OFFICE OF RESEARCH AND DEVELOPMENT
U.S. ENVIRONMENTAL PROTECTION AGENCY
RESEARCH TRIANGLE PARK, NORTH CAROLINA 27711
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DISCLAIMER
This report has been reviewed by the Environmental Sciences Research
Laboratory, U.S. Environmental Protection Agency, and approved for
publication. Approval does not signify that the contents necessarily
reflect the views and policies of the U.S. Environmental Protection Agency,
nor does mention of trade names or commercial products constitute endorse-
ment or recommendation for use.
11
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ABSTRACT
Ambient levels of atmospheric fluorocarbons , halogenated hydrocarbons and
Sp£ are reported for various times and locations in the U.S.A. Requisite
analytical and calibration procedures have been developed. A novel pulsed
flow coulometry gas chromatographic analysis is described. For many com-
pounds, laboratory test-chamber simulations of stratospheric and tropo-
spheric irradiation have been performed. Substrate reactivity and product
profiles are presented, including the effects of the presence of nitrogen
oxides, humidity and a hydrocarbon mixture.
The field studies indicate compounds such as CCl^F, CC12F2' CH^CCl, and
CC14 to be ubiquitous at generally sub ppb levels. Some tropospherically
reactive compounds such as €2014 and C2HC13 are frequently measurable,
while other non-ubiquitous compounds are measurable only where a reason-
able source can be invoked.
The laboratory simulations establish the tropospheric stability of
CC12F2, CH3CC13, CC14 and CC12F-CC1F2. The stratospheric reactivity of
CC14 and CC12F2 is confirmed.
CC14 is a product of C2C14 tropospheric reactions, and should be con-
sidered a secondary anthropogenic pollutant of concern in potential
stratospheric ozone destruction mechanisms. Another product, COC12, is
of potential toxicological concern.
The adventitious labelling of air masses was used to demonstrate urban
ozone transport to rural areas. It is suggested that the controversy
over the origins of non-urban ozone may be resolved by simultaneously
measuring halocarbons.
This report was submitted in fulfillment of Project Number R 800833 by
the Department of Environmental Science, Rutgers, The State University
of New Jersey under the sponsorship of the Environmental Protection
Agency. Work was completed by August 31, 1975.
111
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TABLE OF CONTENTS
Abstract iii
List of Figures viii
List of Tables xvi
List of Terms and Abbreviations xviii
Acknowledgement xix
Sections:
1. Summary and Conclusions 1
1.1 Analytical Procedures 1
1.2 Atmospheric Behavior of Halocarbons 2
2. General Introduction 4
3. Literature Review
3.1 Fll and F12 8
3.1.1 Use and Production 8
3.1.2 Sinks 9
3.1.3 Reactivity Studies 9
3.1.4 Ambient Measurements 10
3.2 Perchloroethylene 10
3.2.1 Use and Production 10
3.2.2 Reactivity Studies 11
3.2.3 Ambient Measurements 12
3.3 Methyl Iodide 12
3.3.1 Sources, Uses and Toxicity 12
3.3.2 Reactivity Studies 13
3.3.3 Ambient Measurements 22
3.4 Carbon Tetrachloride 22
3.4.1 Sources, Uses and Toxicity 22
3.4.2 Reactivity Studies 23
3.4.3 Ambient Measurements 25
3.5 1:1:1 Trichloroethylene 25
3.5.1 Sources, Uses and Toxicity 25
3.5.2 Reactivity Studies 26
3.5.3 Ambient Measurements 26
3.6 Trichloroethylene 27
3.6.1 Uses, Production and Toxicity 27
3.6.2 Reactivity Studies 28
3.6.3 Ambient Measurements 29
3.7 Ethylene Dichloride 30
3.7.1 Uses, Production and Toxicity 30
3.7.2 Reactivity Studies 31
3.7,3 Ambient Measurements 31
3.8 Vinyl Chloride 31
3.8.1 Sources, Uses and Toxicity 31
3.8,2 Reactivity Studies 35
3.8.,3 Ambient Measurements 36
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3.9 Dichloromethane
3.9.1 Sources, Uses and Toxicity 36
3.9,2 Reactivity Studies 37
3.9.3 Ambient Measurements 38
3.10 Polychlorinated Biphenyls
3.10.1 Sources, Uses and Toxicity 38
3.10,2 Reactivity Studies 40
3.10.3 Ambient Measurements 52
3.11 Review of G.C. Analytical Procedures for
Halocarbons 54
4. Experimental
4.1 Reactors 55
4.1.1 Teflon bags and Myler bags 55
4.1.2 Pyrex Reactor 55
4.1.3 Quartz Reactor 55
4.1.4 Light Intensity Measurements 57
4.2 Preparation of Samples for Irradiation 57
4.2.1 Materials 57
4.2,2 Humidity 58
4.2.3 Temperature 58
4,2.4 Chamber Filling Procedure 58
4.3 Permeation Tubes 59
4.4 Laboratory Analyses 59
4.4.1 Gas Chromatograph 59
4.4.2 Nitrogen Oxides 61
4.4.3 Ozone 61
4.4.4 Light Scattering Aerosol 62
4.4.5 Phenols 62
4.4.6 Aldehydes 62
4.4.7 Phosgene 62
4.4.8 Chloride and Hydrogen Ion 62
4.5 Field Analysis - Mobile Laboratory 64
5. Gas Chromatographic Analytical Procedures for Trace
Levels of Halocarbons and SFfc
5.1 Column Evaluation 66
5.2 Gas Phase Coulometry 79
5.2.1 Theory 80
5.2.2 Results 81
5.2.3 Discussion 81
5.3 Pulse Flow Coulometry 92
5.3.1 Theory 93
5.3.2 Results 94
5,3.3 Discussion 95
6. Ambient Measurements
6.1 Results 106
6.2 Discussion 137
vi
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6.2.1 "Ubiquitous" Halocarbons 137
6,2.2 "Non-Ubiquitous" Halocarbons 237
6.2.3 Ambient Variability of Halocarbons 242
6.2.4 Ambient Halocarbons as Tracer for Large
Air Masses 245
6.2.5 Seagirt Monitoring - General comments 245
6.2.6 New York City Monitoring - General comments 245
6.2.7 Sandy Hook Monitoring - General comments 248
6,2.8 Wilmington, Delaware Monitoring - General
comments 248
6.2.9 Baltimore, Maryland Monitoring - General
comments 248
6.2.10 Wilmington, Ohio Monitoring - General
comments 249
6.2.11 Whiteface Mountain Monitoring - General
comments 249
6.2.12 Bayonne, New Jersey Monitoring - General
comments 249
7. Photochemical Reactivity Studies
7.1 F-ll, F-12 and F113 251
7.1.1 Results 251
7.1.2 Discussion 251
7.2 Perchloroethylene 255
7.2.1 Results 255
7.2.2 Discussion 264
7.3 Methyl Iodide 271
7.3.1 Results 271
7.3.2 Discussion 274
7.4 Carbon Tetrachloride 274
7.4.1 Results 274
7.4.2 Discussion 281
7.5 1:1:1 Trichloroethylene 281
7.5.1 Results 281
7.5.2 Discussion 281
7.6 Trichloroethylene 281
7.6.1 Results 281
7.6.2 Discussion 288
7.7 Ethylene Dichloride 290
7.7.1 Results 290
7.7.2 Discussion 290
7.8 Vinyl Chloride 290
7.8.1 Results and Discussion 290
7.9 Dichloromethane 297
7.10 Polychlorinated Biphenyls 297
7.10.1 Results and Discussion 297
List of Publications and Theses ' 325
References 326
vii
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LIST OF FIGURES
Page
1 U.S. Sales of Halocarbons 1967 - 1973 14
2 Absorption Spectrum for CHsI 15
3 Ultraviolet Absorption Spectra of Some
Chlorinated Biphenyls 41
4 Ultraviolet Absorption Spectra of Two
Monochlorobiphenyls 42
5 Comparative Spectral Distributions: Irradiation
Chember vs. Tropospheric Radiation 56
6 Gas Chromato graphic Electron Capture Calibration
Curve for Phosgene 63
7 Typical Chloride Ion Electrode Calibration Curve 65
8 Chromatograms of Ambient (A,B) and Synthetic (C,D)
Mixtures Obtained with Porapak Q at Optimum
Temperatures 68
9 Chromato gram of an Ambient Air Sample on
Chromosil 310 69
10 Chromato gram of an Ambient Air Sample (A) and a
Standard Vinyl Chloride Mixture in Air (B) on
Carbowax 1500 70
11 Chromatogram on an Ambient Air Sample on S.E. 30 71
12 Chromatogram of an Ambient Air Sample on D.C. 710 72
13 Chromatogram of an Ambient Air Sample on Carbowax 400 73
14 Dual Trace Coulometric Chromatogram of an Ambient Air
Sample on a D.C. 200 Column 76
15 Dual Trace Coulometric Chromatogram of an Irradiated
C2C14/N02 Mixture Showing Phosgene Synthesis. Didecyl
Phthalate on 100/120 Mesh Chromosorb P Column 77
16 Peak Height vs. Concentration for CCU and CC13F 84
17 Effect of Flow Rate on lonization Efficiencies 85
18 Effect of Concentration of CC14 and CC13F on the
Coulometric Response at Fixed Efficiency 86
viii
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Page
19 Convergence of Coulometric Measurements to
Actual Concentration as a Function of
lonization Efficiencies 87
20 Typical Chromatogram of New Brunswick,
New Jersey Air Sample 88
21 Chromatogram of Trans-CHC1=CHC1, CH2=CCl2 Showing
Greater than Coulometric Response 89
22 Effect of Trans-CMC1=CHC1 Concentration on the
Greater than Coulometric Reponse 90
23 Dual E.C. Chromatogram of Standard Mixtures of
CC13F, CC14 and CC12F CC1F2, Respectively in Air 96
24 Weight Loss of Phosgene Permeation Tube with Time 97
25 Dual E.C. Chromatogram of Replicate Injections of
a Standard Phosgene Mixture 98
26 Semi-Log Plot of Ce versus 1/Fo for Several
Standard Concentrations of Phosgene 99
27 Variation of lonization Efficiency of Phosgene
with Flow Rate 100
28 Variation of lonization Efficiency of Phosgene with
Input Sample Concentration 101
29 Phosgene Decay in a Conditioned Glass Vessel in
Presence and Absence of Water Vapor, and in Contact
with a Ground Surface 102
30 Monitoring Locations 107
31 § Seagirt, New Jersey 164
31A 6/18/74 Halocarbon Levels 165
32 § Seagirt, New Jersey 166
32A 6/19/74 Halocarbon Levels 167
33 § New York City 168
33A 6/27/74 Halocarbon Levels 169
34 & New York City 170
34A 6/28/74 Halocarbon Levels 171
IX
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Page
35 S Sandy Hook, New Jersey 172
35A 7/2/74 Halocarbon Levels 173
36 $ Sandy Hook, New Jersey 174
36A 7/3/74 Halocarbon Levels 175
37 $ Sandy Hook, New Jersey 176
37A 7/4/74 Halocarbon Levels 177
38 § Sandy Hook, New Jersey 178
38A 7/5/74 Halocarbon Levels 179
39 5 Wilmington, Delaware 18°
39A 7/8/74 Halocarbon Levels 181
40 £ Wilmington, Delaware 182
40A 7/9/74 Halocarbon Levels 183
41 § Wilmington, Delaware 184
41A 7/10/74 Halocarbon Levels 185
42 § Baltimore, Maryland 186
42A 7/11/74 Halocarbon Levels 187
43 § Baltimore, Maryland 188
43A 7/12/74 Halocarbon Levels 189
44 § Wilmington, Ohio 19°
44A 7/21/74 Halocarbon Levels 191
45 § Wilmington, Ohio
45A 7/22/74 Halocarbon Levels
46 § Wilmington, Ohio
46A 7/23/74 Halocarbon Levels
47 § Wilmington, Ohio 196
47A 7/24/74 Halocarbon Levels 197
i no
48 § Wilmington, Ohio
48A 7/25/74 Halocarbon Levels 199
49 £ Whiteface Mountain, New York
49A 9/16/74 Halocarbon Levels 201
50 § Whiteface Mountain, New York
50A 9/17/74 Halocarbon Levels
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Page
51 £ Whiteface Mountain, New York 204
51A 9/18/74 Halocarbon Levels 205
52 § Whiteface Mountain, New York 206
52A 9/19/74 Halocarbon Levels 207
53 Seagirt, New Jersey 6/18/74 Ambient Data 208
54 Seagirt, New Jersey 6/19/74 Ambient Data 209
55 New York City 6/27/74 Ambient Data 210
56 New York City 6/28/74 Ambient Data 211
57 Sandy Hook, New Jersey 7/2/74 Ambient Data 212
58 Sandy Hook, New Jersey 7/3/74 Ambient Data 213
59 Sandy Hook, New Jersey 7/4/74 Ambient Data 214
60 Sandy Hook, New Jersey 7/5/74 Ambient Data 215
61 Wilmington, Delaware 7/8/74 Ambient Data 216
62 Wilmington, Delaware 7/9/74 Ambient Data 217
63 Wilmington, Delaware 7/10/74 Ambient Data 218
64 Baltimore, Maryland 7/11/74 Ambient Data 219
65 Baltimore, Maryland 7/12/74 Ambient Data 22°
66 Wilmington, Ohio 7/19/74 Ambient Data 221
67 Wilmington, Ohio 7/21/74 Ambient Data 222
68 Wilmington, Ohio 7/22/74 Ambient Data 223
69 Wilmington, Ohio 7/23/74 Ambient Data 224
70 Wilmington, Ohio 7/24/74 Ambient Data 225
71 Wilmington, Ohio 7/25/74 Ambient Data 226
72 Wilmington, Ohio 7/26/74 Ambient Data 227
73 Wilmington, Ohio 7/27/74 Ambient Data 228
xi
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Page
74 Whiteface Mountain, New York
9/16/74 Ambient Data 229
75 Whiteface Mountain, New York
9/17/74 Ambient Data 230
76 Whiteface Mountain, New York
9/18/74 Ambient Data 231
77 Whiteface Mountain, New York
9/19/74 Ambient Data 232
78 Seagirt, New Jersey
6/19/74 Atmospheric Measurements 233
79 New York City 6/27/74 Atmospheric Measurements 234
80 Sandy Hook, New Jersey
7/2/74 Atmospheric Measurements 235
81 Baltimore, Maryland
7/12/74 Atmospheric Measurements 236
82 Halocarbon (In Air) +50 pphm N02 + 50% RH +
hv in a 72 liter Pyrex Glass Reactor, CC12F2 and CC13F 238
83 Halocarbon (In Air) + 50 pphm N02 + 50% RH +
hv in a 200 liter Mylar Bag, CH3CC13 and CC14 239
84 C2C14 (In Air) + 50 pphm N02 + 50% RH -t- hv
in a 200 liter Teflon Bag 241
85 CHsI (10 ppb in Air) + 50 pphn N02 + 50% RH
+ hv in a 200 liter Teflon Bag 243
86 Halocarbon (In Air) + 50 pphm N02 + 50% RH
+ hv in a 72 liter Pyrex Glass Reactor, CC12-CCC1F2 244
87 Whiteface Mountain, New York 9/17/74 246
88 New York City (45th St. § Lexington Ave.) 6/27/74 247
89 CClsF plus N02 and Hydrocarbon Irradiated in Ultra
Zero Air in a Mylar Reaction Chamber
90 Freon 11, Freon 12 plus N02 Irradiated in Ultra Zero
Air in a 72 liter Pyrex Glass Reactor, Replenishing
N02
xii
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Page
o
91 Hydrocarbon decay in 1 liter U.V. (>2200A)
Reactor, CC12F-C1F2, C1C13F, CC14 254
92 £2014 Irradiated in Ultra Zero Air 257
93 C2C14 Plus Hydrocarbon Irradiated in Ultra Zero Air 258
94 C2Cl4 Irradiated in N2 259
95 C2d4 Plus N02 and Hydrocarbon Irradiated in
Ultra Zero Air - Run 1 260
96 C2C1^ Plus N02 and Hydrocarbon Irradiated in
Ultra Zero Air - Run 2 261
97 C2Cl4 Plus N02 and Hydrocarbon Irradiated in
Ultra Zero Air - Run 3 262
98 C2C14 Plus N02 and Hydrocarbon Irradiated in
Ultra Zero Air - Run 4 263
99 Ozone Dark Reactions. Ozone Plus C2C14 267
100 Duplicate Runs 268
101 N02 Photolysis in Air, 50% RH in Teflon Bag 272
102 CH3I + 03 (dark) 273
103 CH3I Photolysis in N2 275
104 CH3I Photolysis in N2, Semi-Log Plot 276
105 CH3I Photolysis in Air, 50% RH 277
106 CH3I Photolysis in Air, 50% RH, Semi-Log Plot 278
107 CH3I + N02 Photolyzed in Air, 50% RH, Semi -Log Plot 279
108 CH3I + N02 Photolyzed in Air, 50% RH, 1 ppra
Hydrocarbon Present 280
109 CC14 and CH3CC13 Irradiated in a 144 liter
Mylar Bag with 1 ppm Halocarbon, 50% RH and
0,5 ppm N02 Added Daily 282
110 Trichloroethylene + 03, dark reaction 283
111 Trichloroethylene in N2 Irradiated in a
200 liter Teflon Bag 284
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Page
112 Trichloroethylene in Ultra Zero Air + 50% RH
Irradiated in a 200 liter Teflon Bag 285
113 Trichloroethylene + N02 + 50% RH + Air
Irradiated in a 200 liter Teflon Bag 286
114 Trichloroethylene + N02 Fuel A + 50% RH
Irradiated in a 200 liter Teflon Bag 287
115 Dichloroethylene + NCb + Fuel A + 50% RH
Irradiated in a 200 liter Teflon Bag
Replenishing N02 Daily 291
116 Photolysis of Vinyl Chloride in Dry Nitrogen,
in Air, 50% RH 293
117 Photolysis of Vinyl Chloride in Air, 50% RH
N02 Present 294
118 Photolysis of Vinyl Chloride in Air, N02 and
Hydrocarbon Present, 50% RH 295
119 "Dark" Reaction of Vinyl Chloride and Ozone in
Air, 50% RH 296
120 Ortho-Chlorobiphenyl Stability in 250 liter Teflon
Bags in the Absence of Ultraviolet Radiation 299
121 Ortho-Chlorobiphenyl in Nitrogen Irradiated with
Simulated Tropospheric Radiation (linear plot)
122 Ortho-Chlorobiphenyl in Nitrogen Irradiated with
Simulated Tropospheric Radiation (Log Concentration
vs. Time) 301
123 Ortho-Chlorobiphenyl in Nitrogen Irradiated with
Fourfold 310 Nanometer Intensity Source (Linear Plot)
124 Ortho-Chlorobiphenyl in Nitrogen Irradiated with
Fourfold 310 Nanometer Intensity Source (Log
Concentration vs. Time) 303
125 Ortho-Chlorobiphenyl in Air Irradiated with
Simulated Tropospheric Radiation (Linear Plot) 305
126 Ortho-Chlorobiphenyl in Air Irradiated with
Simulated Tropospheric Radiation (Log
Concentration vs. Time) 306
xiv
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Page
127 Soluble Chloride and Hydrogen Ion Production
in the Nitrogen/Fourfold 310 Nanometer Intensity
Irradiation (Log Concentration vs. Time) 308
128 Soluble Chloride and Hydrogen Ion Production
in the Nitrogen/Simulated Tropospheric Irradiation
(Log Concentration vs. Time) 309
129 Soluble Chloride and Hydrogen Ion Production
in the Air/Simulated Tropospheric Irradiation
(Log Concentration vs. Time) 310
130 Chloride Ion: Ortho-Chlorobiphenyl Loss vs. Time
in the Nitrogen/Fourfold 310 Nanometer Intensity
Irradiation 311
131 Chloride Ion: Ortho-Chlorobiphenyl Loss vs. Time
in the Nitrogen/Simulated Tropospheric Irradiation 312
132 Chloride Ion: Ortho-Chlorobiphenyl Loss vs. Time
in the Air/Simulated Tropospheric Irradiation 313
133 Soluble Hydrogen Ion Formation: Ortho-Chloro-
biphenyl Loss vs. Time in the Nitrogen/Fourfold
310 Nanometer Intensity Irradiation 314
134 Soluble Hydrogen Ion: Ortho-Chlorobiphenyl Loss
vs. Time in Nitrogen/Simulated Tropospheric
Irradiation 315
135 Soluble Hydrogen Ion: Ortho-Chlorobiphenyl Loss vs.
Time in the Air/Simulated Tropospheric
Irradiation 316
136 Phenolics Formation in the Nitrogen/Fourfold 310
Nanometer Intensity Irradiation 318
137 Phenolics Formation in the Nitrogen/Simulated
Tropospheric Irradiation 319
138 Phenolics Formation in the Air/Simulated
Tropospheric Irradiation 321
139 Light Scattering Aerosol Formation in the Nitrogen/
Simulated Tropospheric Irradiation 322
140 Light Scattering Aerosol Formation in the Air/
Simulated Tropospheric Irradiation 323
xv
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List of Tables
Tables Page
1 U.S. Production of F-12 and F-ll
(U.S. Tariff Commission Reports)
(units of 1000 pounds). 8
2 Synthetic Organic Chemicals, U.S.
Production and Sales (U.S. Tariff
Commission Reports) (Units of
1000 pounds) 11
3A-C Previous PCB Studies - Summary of
Conditions and Results 45
4 Permeation Tube Data 60
5 Useful Columns and Temperatures for
Ambient Halocarbon Analysis 67
6A&B Comparison of Coulometrically Calcu-
lated concentrations with Calibra-
tion Standards 78
7 lonization Efficiencies 82
8 Ambient Concentrations of Coulo-
metrically Determined Compounds in
the New Brunswick, N.J. Area 83
9 Effect of Flow Rate on Detector
Responses and Efficiencies for CClsF
and CH31 83
10 Comparison of the Coulometrically
Determined Concentrations of CC^F,
CCl4 and CCl2F-CClF2with Permeation
Standards 103
11 Comparison of Standard Phosgene
Concentrations with PFC - Calculated
Concentrations 104
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List of Tables
Table
12A-S
13A&B
14A&B
15A-D
16A&B
17
18A-W
19A&B
20
21
Halocarbons and SFg Ambient Levels
at Various Times and Locations:
Complete Data
Maximum and Minimum Halocarbon Levels
At Various Locations in the United
States. Concentrations (ppb)
Typical Levels of Halocarbon
Halocarbon and SF5 Ambient Levels
at Various Times and Locations -
Daily Mean and Monitoring Period
Mean Values (Parenthesis indicate
number of observations at minimal
detectable values)
Percentages of Analyses in which
Each Compound was Detectable
Halocarbons and SFg in Order of
Ambient Variability - Expressed
As Weighted Average S.D. for
Each compound
Aerometric Parameters at Various
Times and Locations
Scale Factors Used in Figures
31-77
Perchloroethylene Systems inves-
tigated
Major Products of Photodissocia-
tion C2C14 in Air
Page
108
127
129
131
135
138
139
162
265
266
xvi i
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LIST OF TERMS AND ABBREVIATIONS
Fll Freon 11,
F12 Freon 12, CC12F2
F113 Freon 113, CC12F-CC1F2
PCB's Polychlorinated biphenyls
G.C. Gas Chromatography
E.G. Electron Capture
P.F.C. Pulse flow coulometry
•C, Boundary condition
STUV Simulated Tropospheric Ultraviolet Radiation
OCB Ortho-chlorobiphenyl
FEP Fluorinated ethylene propylene copolymer
PAN Peroxyacetyl nitrate
1.1.l.T 1:1:1 Trichloroethane
T.C.E. Trichloroethylene
P.C.E. Perchloroethylene (Tetrachloroethylene)
V.C. Vinyl chloride
F.S.D. Full Scale Deflection
xvi 11
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ACKNOWLEDGEMENT
The following have contributed significantly to this project: M, Antell,
A. Appleby, R. Arnts, S. Bergelson, R, Gumpert, R. Hague, D, Hansen,
J. Kazazis, D. Lillian, L. Lobban, B. Scott, H.B, Singh, and J. Toomey,
Thanks are due to Dr. G. Wolff (Interstate Sanitation Commission, New
York), and Dr. W.N. Stasiuk (New York State Department of Environmental
Conservation) for their help in arranging some of the field studies and
for useful discussion.
xix
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SECTION I
SUMMARY AND CONCLUSIONS
There is much recent concern over the behavior and effects of halocarbons
in the environment. Ambient CC13F and CCl2F2^1'^ ) and, by analogy, other
tropospherically stable compounds are suspect as precursors of strato-
spheric ozone-destroying chlorine atoms. Vinyl chloride has been linked
to industrial angiosarcoma and is possibly mutagenicf^ ). Chloroethylenes
will react to form significant quantities of highly toxic phosgene and
acetyl chlorides under simulated tropospheric conditions^ > J. Chloro-
form and CC14 in Mississippi drinking water have been associated with an
elevated cancer riskC ' ). Concurrent to these findings, the research
effort in ambient halocarbon measurement has justifiably increased signifi^
cantly(8,9,10,11,12,13,14,15,16).The many industrial and domestic uses of
halogenated hydrocarbons and their rather large production figures (12
billion pounds in 1974 for the U.S. alone)^ ) suggest that this accel-
erated research effort will lead to a continual increase in the number
of halocarbons routinely measured in the environment. Indeed, mass spec-
trometric analysis conducted on cryogenically concentrated air samples
several years ago indicated the presence of a wide variety of halogenated
compoundsC 18).
Because of the large number of atmospheric halogenated hydrocarbons of
potential interest, a selective study of a carefully chosen group rep-
resenting a wide spectrum of chemical reactivities and emission patterns
was clearly desirable. This approach would not only provide information
of immediate relevance but also contribute to a data base for future
reference as the number of halocarbons of environmental interest pro-
liferates. Accordingly, this study has been directed towards a) develop-
ment of analytical procedures for the determination of trace levels of
halocarbons and SF^ in the atmosphere, b) monitoring as many of these
compounds as possible at a variety of locations in the United States,
coupled with observation of their photochemical behavior in clean and
"polluted" air under simulated tropospheric (and in some cases strato-
spheric) conditions.
1.1 Analytical Procedures
Gas chromatography, using electron capture (E.G.) detection was used
throughout the study, and a number of different types of column packing
were evaluated. DC 200 has been demonstrated to be an excellent packing
for the separation and E.G. determination of ambient CC13F, CH3I, CC^F-
CC1F2, CHC13, CH3-CC13, CC14, C2HC13 and C2C14 under isothermal (23°C)
conditions. Porapak Q at 80°C, and Chromosil 310 and Carbowax 1500 at
23°C, were found acceptable for ambient CC12F2 analysis. Poropak Q at
room temperature is also a very good column for the separation of SFg
and CBrF3. Carbowax 1500 exhibited excellent resolution of vinyl chloride
at room temperature (23°C). Phosgene was found to suffer unacceptable
irreversible losses in all columns studied except didecyl phthalate. This
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column provided excellent resolution at 23°C and sorption losses were
sufficiently small to allow accurate analysis in the sub-ppb range with
frequent calibrat ion.
The use of purge gas at low concentrations of solute leads to reduced
sensitivity. This is attributable to a decrease in ionization efficiency
with increasing flowrate through the EC detector. The utilization of the
ionization efficiency as a confirmatory aid has been demonstrated and
further applied to confirm the identity of phosgene in a simulated photo-
chemical smog study and, accordingly, the probable tropospheric synthesis
of this compound. Because of its high ionization efficiency (>75%) and
instability, the latter making the preparation of calibration mixtures
extremely difficult, phosgene and similar reactive strong electron absorbers
are excellent candidates for coulometric analysis. Accordingly, for such
compounds a detailed evaluation of gas phase coulometry and analysis of
sorption kinetics was indicated. Gas phase coulometry was evaluated and
used for the measurement of a number of ambient halocarbons since the
method permits operation at maximum sensitivity and is absolute. Ionization
efficiencies were determined for 14 compounds to range from less than 0 to
90%. Reciprocal efficiencies were verified to increase linearly with flow
rate as predicted by a stirred tank reactor model. At intermediate ef-
ficiencies (40 to 60%), coulometric measurements of CCljF and CC14 exhi-
bited a constant error of about 25% over a representative concentration
range. At efficiencies greater than 90%, the percent error in coulometric
analysis was less than 5% compared to primary standards. Compounds iden-
tified and measured coulometrically included CClgF, CH^I, CH^-CCl^, CC1./J,
CHC1=CC12 and CC12=CC12.
For the absolute analysis of reactive electron-absorbing compounds which
undergo decomposition in a G.C. column, a novel modification of gas phase
coulometry, termed Pulse Flow Coulometry (PFC) was developed and used
successfully. Comparing this technique against permeation tube standards,
an error of less than 15% was easily achieved when ionization efficiencies
were greater than 75%. At ionization efficiencies greater than 85%, this
error could be reduced to 4%. The electron capture detector response to
phosgene was comparable to that to carbon tetrachloride, one of the
strongest known electron absorbers. It was demonstrated that ppb mixtures
of phosgene in air are quite stable in the presence of moisture and undergo
rapid heterogeneous decay on surfaces. PFC should be useful for the
analysis of any unstable electron absorber (e.g. peroxyacetyl nitrate [PAN],
chlorine dioxide, trichloroacetyl chloride) in air and offers a wide scope
of application in air pollution and industrial hygiene.
1.2 Atmospheric Behavior of Halocarbons
The halogenated compounds studied are primarily anthropogenic and exhibit
wide fluctuations in concentrations as a function of location and meteoro-
logical factors. However, typical halocarbon ranges at different locations
can be inferred from the data base presented here. CClsF, CC12F2, CH3CC13
and CC14 were found to be ubiquitous atmopsheric constituents. Their
concentrations in the rural areas were typically in the sub-ppb range and
-------
urban concentrations were generally much higher. The first three compounds
are of known anthropogenic origin but the CC14 atmospheric budget is in-
consistent with direct anthropogenic emissions,
Our studies support the possibility of the formation of some CC14 in the
troposphere by solar-induced photochemical reactions of chlorinated
alkenes. Such a process is of major importance in the consideration of
ozone destruction by chlorine atoms, since our laboratory simulation
studies indicate that CC14, like the fluorocarbons, is tropospherically
stable but would break up under stratospheric conditions. Chlorinated
alkenes should therefore be included in any evaluation of possible ozone
destruction. The ubiquitous compounds and CC12F-CC1F2 are tropospherically
stable and will experience their ultimate fate in the stratosphere. Among
the tropospherically reactive halocarbons, C2C14 and C2HC13 should lead to
the formation of large quantities of phosgene in the atmosphere which may
have a potentially significantly impact. Methyl iodide was measurable
generally only near the ocean, supporting its origin there. SF6 and
vinyl chloride were usually measurable only where a reasonable nearby
source could be involved.
It has been demonstrated that halocarbons can be used as excellent tracers
of urban transport. Additionally, studies of the temporal and spatial
distributions of halocarbons in non-urban areas should help to differen-
tiate between the natural synthesis, urban transport and stratospheric
injection of ozone.
Because of the possible stratospheric ozone destruction by stable halo-
carbons, the toxicity of phosgene, the indicated carcinogenicity of vinyl
chloride and its structural similarity to other ambient chloro-ethylenes,
and the unknown long-term effects of halogenated compounds and their
reactive products, further research in this area is clearly warranted.
-------
SECTION II
GENERAL INTRODUCTION
This project had as its inception a literature search which was initiated
in the Spring of 1971 to estimate worldwide production of industrial
chemicals and the percentages of them emitted into the atmosphere and to
consider the tropospheric removal mechanisms and fates of these materials.
As a result of this search it was determined that in terms of the quantities
of materials estimated to be emitted on the one hand and the possible
environmental impact of these primary pollutants or their degradation
products on the other hand, continuation of this effort in the direction
of ambient monitoring and of tropospheric and stratospheric simulation
experiments was justified. It was especially clear from the initial
literature search that there was a paucity of data on the environmental
fates of halocarbons in general and fluorocarbons in particular. This
followed and was understandable from the almost complete lack of per-
tinent information on the atmospheric concentrations of these materials
and on their sinks and on the identity and kinetics of their reactants,
intermediates and products of photochemical degradation. As a result, a
proposal was submitted to E.P.A. in 1972 to provide requisite information
for subsequent meaningful toxicological and environmental impact studies.
The initial proposal listed 19 chemicals, most of which were halogenated
hydrocarbons. Production figures taken from the pertinent U.S. Tariff
Commission Reports were listed, and from these, projections on worldwide
sales in the categories of aerosol propellants, refrigerants, plastics,
solvents, blowing agents and miscellaneous were developed. It was shown
that 99% of the world's fluorocarbon sales were accounted for by only 5
Freons and plastics, and that major fluorocarbon usage was in aerosol
propellants.
Various possible sinks for the fluorocarbons were considered, viz: dis-
solution in the ocean, washout by rain, escape into the stratosphere,
microbial utilization, absorption at the earth's surface, air breathers
and plants and tropospheric photochemical reactions. Dissolution in the
oceans was shown to be insignificant from a simple consideration of
Henry's Law. Washout by rain was similarly insignificant. Microbial
utilization and absorption at the earth's surface were similarly ruled
out, and later a detailed analysis of mechanical and biological air
breathers showed them to be insignificant sinks. The possibility of
sinks involving stratospheric and tropospheric photochemical reactions
and plants was left open. Indeed the only hint at the relative impor-
tance of these possibilities was provided by a few isolated measurements
of atmospheric concentrations of F-ll and SFg by Lovelockt^. He repor-
ted the F-ll background concentration as 10~*1 v/v. Computing the total
worldwide emissions of F-ll from 1958 to the time of Lovelock's measure-
ments gave a total worldwide tropospheric emission of F-ll during the
period of 1958 to July 1970 as 3376xl06 pounds. Comparing this with the
-------
actual "tropospheric loading" calculated from Lovelock"s background
measurements of 553x10^ pounds it was concluded that important sinks for
fluorocarbons existed. The known tropospheric/stratospheric turnover
rate could only account for a fraction of this deficit strongly supporting
the existence of tropospheric photochemical sinks involving reactive
hydrocarbons, oxides of nitrogen, 63 and their associated intermediates.
Also, a more extensive study was indicated because of the significance
of possible build-up of these compounds towards such factors as toxico-
logical effects, global greenhouse effect, and their potential use as
tracers of air masses. Accordingly the proposed research protocol
included tropospheric photochemical reactivity studies of saturated halo-
carbons as well as those unsaturated halocarbons emitted in significant
quantities and expected to be reactive. Because, despite the available
ambient measurements, all theoretical considerations would argue against
reaction of the saturated halocarbons in the troposphere, the original
proposal also included an elaborate plan to study stratospheric reac-
tivity of many of these compounds, with emphasis on the fluorocarbons.
Indeed, literature existed that the halocarbons would undergo photolytic
decay at wavelengths corresponding to stratospheric conditions. E.P.A.
did not fund this stratospheric halogenated hydrocarbon aspect of the
proposal. Some modest studies were however carried out and are reported
here. The role of plants as possible sinks was not considered in this
study.
Also included with the fluorocarbons for study were such possibly car-
cinogenic compounds as vinyl chloride, CC14 and trichloroethylene and
the teratogenic P.C.B.'s. Product identification was also considered to
be an important feature because of the possibly significant toxicity of
such expected products as phosgene and the chloroacetyl chlorides. The
role, if any, that reactive halogenated compounds play in the classical ,
photochemical smog cycle was also a feature of the work.
In order to approach the assessment of the fates of these compounds the
project involved three phases running essentially concurrently. (1) A
continuation of the literature survey to upgrade our information on
emissions data and possible sinks. (2) A program of ambient measure-
ments to provide tropospheric concentrations of the compounds and to
estimate their atmospheric budget and removal mechanisms. (3) Laboratory
studies in which controlled simulation of tropospheric and stratospheric
photochemical conditions were used. It was necessary to develop analy-
tical procedures and techniques usable at the very low concentrations
(10~9 to 10~12 v/v) anticipated, as a prelude to the required program of
ambient measurements. Also, significant research and development effort
had to be expended on the preparation, irradiation and sampling of gas
mixtures in a fashion consistent with optimum simulation of ambient
conditions.
The arrangement of this report is as follows. Following this General
Introduction is a section on Literature Review, with subsections on the
major compounds studied. Then follows the Experimental Section which
-------
includes all the standard analytical procedures used in this work. A
considerable portion of this project involved development of gas chroma-
tographic procedures and columns for the analysis of very low levels of
halocarbons and related compounds. While a brief account of the gas
chromatographic instrumentation used appears in the Experiment Section,
a complete account of this analytical development is given in the next
section, "Gas Chromatographic Analytical Procedures for Trace Levels of
Halocarbons and SFg."
Section VI presents the results and discussion of our ambient measure-
ments, with subsections devoted to "ubiquitous" and "non-ubiquitous"
halocarbons, their variability, and general comments pertaining to each
of the locations studied. Section VII describes the photochemical
reactivity studies with subsections for each compound.
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SECTION III
LITERATURE REVIEW
INTRODUCTION
An important aspect of this work was a continuing review of available
literature on the sources, reactions, fates and levels of these compounds.
Presented in this section are references to work of especial significance
to this project.
As mentioned in the General Introduction, the earliest report on background
concentrations of any of these compounds was that of Lovelock(*0 . He re-
ported CC13F and SFg values in southwest Ireland of 10~n and 2.9xlO~14
(v/v), respectively, as representative of northern hemispheric background
concentrations. Air passing over the continent from an easterly direction
exhibited corresponding concentrations approximately an order of magni-
tude higher for both. Atmospheric concentrations ranging from 0.05 to
0.17 ppb CC13F in Bowerchalke, England, were also reported by Lovelock'•-'
for the period October 1970 to October 1971. No significant seasonal
variation of CC13F was observed. Lovelock et al.( ^also reported mean
aerial concentrations of 0.049 CC13F, 0.001 CH3I and 0.071 CC14 ppb (v/v)
over the north and south Atlantic Ocean. Hester et al.(") reported CC13F
and CC12F2 levels at various indoor and outdoor locations in the Los
Angeles Basin. Ambient CC12F2 and CC13F levels were in the range of
0.1 to 1.25 ppb and 0.05 and 147 ppb respectively. Simmonds et al.(16)
in a three-day study reported levels of CC13F, CC14, CH3CC13 and C2C14
in the Los Angeles Basin. The reported concentration (ppb) ranges were
0.11-2.2 for CC13F, 0.1-1.63 for CC14, 0.01-2.3 for CH3CC13 and 0.1-4.2
for C2C14.
Wilkness, Lamontagne, Larsen and Swinnerton( 'reported CC14 and F-ll
concentrations over the ocean of about 80 and 60 ppb respectively. Re-
cently Lovelock(19J reported concentrations of F-ll, F-12, CHCls, HIT,
CC14, TCE, and PCE in western Ireland (June, July 1974) and over the
north Atlantic Ocean (October 1973). Except for TCE, the values in the
two locations were approximately the same, namely:
F-12 about 100-115 ppt.
F-ll about 80-90 ppt.
CHC13 about 19-27 ppt.
HIT about 65-75 ppt.
CC14 about 111-138 ppt.
TCE <5 in north Atlantic
15 in Ireland
PCE 21-28
(It should be noted that the F-ll values are about an order of
magnitude greater than his earlier report.)
He also reported vertical profiles of F-ll and CC14 over central England
(June 6, 1974), demonstrating their presence in the tropopause and lower
7
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stratosphere. Saltzman et al.*' irradiated mixtures of reactive atmos-
pheric pollutants, CBrFs, SFg and C4Fg with simulated sunlight and
attributed the observed loss of less than 2% per day for the above com-
I O"^
pounds to permeation through the walls of the film bags. Hester et al. *• J
provided support for the short-term tropospheric stabilities of CC12F2
and CC13F based on smog simulation studies. Wilson^ •* reported the photo-
reactivity of trichloroethylene and its possible involvement in the photo-
chemical smog phenomenon. Molina and Rowland ^ ' and Cicerone et al.^ '
have recently suggested the possible destruction of the stratospheric
ozone lyaer by chlorine-atom-propagated chain reactions initiated by the
photolysis of CCljF and CC12F2- Experimental smog chamber data supporting
the assumed long-term tropospheric stability of these compounds, however,
is insufficient. No data at all exist on the tropospheric stability of
many other chlorinated compounds present in the troposphere at comparable
concentrations to CC^F and CC12F2-
SPECIFIC COMPOUNDS
3.1 F-ll and F-12
3.1.1 Use and Production -
Freon 11 and Freon 12 are used widely in household and industrial environ-
ment with extensive application as refrigerants, propellants for per-
fumery and medicinal products, insecticides and fire repellants. They
are being used more and more in space technology and their worldwide
uses are even expected co increase as underdeveloped countries become
more and more industrialized. Dupont (1972) estimated that 99% of the
world's fluorocarbon sales are accounted for by Freon 11, 12, 22, 113 and
114. Freon 11, 12 and 134 are the major fluorocarbons used for aerosol
propellants.
The use of the Freon 11 and Freon 12 in household and industrial appli-
cation is increasing at a significant rate as seen from U.S. Production
Data (1965-1973).
Table 1. U.S. PRODUCTION OF F-12 AND F-ll
(U.S. Tariff Commission Reports)
(Units of 1000 pounds)
Year Freon-12 Freon-11
1965 271,408 170,461
1966 286,326 170,350
1967 309,668 182,216
1968 325,625 204,418
1969 363,658 238,518
1970 375,406 244,472
1971 389,580 257,899
1972 439,224 299,583
1973 488,831 333,773
8
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The 1972 world production rates for Freon 11 and Freon 12 are about 0.3
and 0.5 million tons per year respectively^ ) ancj are steadily increasing.
By the nature of the use of these compounds they are emitted to the tro-
posphere in significant quantities.
3.1.2 Sinks -
An all inclusive list of possible sinks for any tropospheric constituent
would include the following:
(1) Dissolution into the ocean according to Henry's Law followed by
hydrolysis.
(2) Washout by rain.
(3) Escape through the tropopause gaps to the stratosphere.
(4) Microbiological utilization at earth surface.
(5) Absorption at earth surface.
(6) Sorption and/or utilization, and/or degradation by plant life and
biological and mechanical air breathers.
(7) Stratospheric and tropospheric photochemical and/or thermal reactions.
Careful analysis indicates that due to such properties as chemical inert-
ness, low solubility in water, slow hydrolysis rate, lack of data to
indicate biodegradability and low partial pressure of fluorocarbons, the
aforementioned sinks numbered 1, 2, 4 and 6 are very insignificant in
terms of the total amount of Freon 11 and Freon 12 they would remove
from the troposphere. Freon 11 and Freon 12 are not very soluble in
water and therefore not removed by rainouts in the troposphere^ J. The
relative insolubility in water together with the chemical stability,
particularly towards hydrolysis( •> indicates that these molecules will
not be rapidly removed by dissolution in the ocean, and measurements
made indicate equilibrium between the ocean surface and air. Therefore
a major sink other than photolysis cannot be inferred*- '. Based on the
evidence presented in this discussion tropospheric and thermal reactions
or stratospheric flux must be considered as the significant removal
mechanism.
3.1.5 Reactivity Studies -
Hester^ -* reported long term photolysis studies on Freon 11 and Freon 12
in which 76 ppb Freon 12 and 2.3 ppb Freon 12 were irradiation in Los
Angeles air. However, after almost 2 months photolysis, no observation
of a chemical reaction of either Freon 11 or Freon 12 was discovered. A
second scheme of experimentation was pursued in which a photolysis study
was run analagous to that previously stated except that 0.1 ppm N02 was
added to the system. The photolysis was for 7 hours and the experimenters
noted no indication of any chemical reaction, as indicated by no de-
crease in either fluorocarbon concentration. The photolysis studies were
done in a 20 liter reaction flask surrounded by 11 "Black Light" Type U.V.
lamps.
-------
5.1.4 Ambient Measurements -
Ambient measurements of Freon 11 and Freon 12 have indicated their
presence in the atmosphere. LovelockC^) reported a measured concentration
exceeding 10~H by volume over Ireland. Measurements of CCl^Y have also
been made in several parts of the world. Lovelock et al. @Q) reported
ambient atmospheric measured concentration of Freon 11 for the Los Angeles
Basin, Europe, and other parts of the Atlantic. Based on these measure-
ments they estimated a background concentration of 1x10"^ v/v. Wilkniss
et al. U-3) measured trace gases in the atmosphere over the north and south
Pacific Ocean. Goldberg et al. (•*-•>) reported average atmospheric concen-
trations of 2.9xlO~10 and 3.2x10-9 for CC1$F and CCl2?2 respectively.
Simmonds et al. 0-"J measured the concentration of four halogenated hydro-
carbons throughout the air over the Los Angeles Basin in which they
reported average concentrations of .55 ppb for CCl^F. Based on ambient
atmospheric measurements of Freon 11 and Freon 12 in the greater Los
Angeles Basin, it was found(9) that concentrations of Freon 11 and
Freon 12 in homes range from .22 to .12 ppb for Freon 11 and 0.3 to
510 ppb for Freon 12.
It is of particular interest to note that this study indicated that the
levels of Freon 11 and Freon 12 were significantly higher in homes than
in ambient air samples.
5.2 Perchloroethylene
5.2.1 Use and Production -
In January 1963 there were an estimated 36.500 dry cleaning plants in
the United States resulting in approximately one plant for every 2500
people in urban areas and one for each 500 people in the country in
general( ). There were also 7300 self service "coin operated" dry
cleaning stores containing about 50,000 machines. It was also estimated
at that time that the dry cleaning industry alone consumed 25 million
gallons of perchloroethylene per year. It is of importance to note that
because of zoning laws against petroleum solvents, the number of plants
utilizing perchloroethylene is increasing at a rapid rate, and that for
reasons of convenience dry cleaning plants are located at or near centers
of high population density. Estimates indicated that in 1968 the dry
cleaning industry accounted for 75% of the reported sales of perchloro-
ethylene^"J. The remainder probably also results in significant emis-
sion of the compound in urban atmospheres. The following table indicated
the increase in the production of perchloroethylene from 1965 to 1972.
10
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Table 2. SYNTHETIC ORGANIC CHEMICALS, U.S. PRODUCTION AND SALES
(U.S. Tariff Commission Reports)
(Units of 1000 pounds)
1965
1966
1967
1968
429,354
462,678
532,980
636,484
1969
1970
1971
1972
635,251
706,896
704,747
734,216
The above report indicated an increase of approximately 70% in the U.S.
production of perchloroethylene over this period.
3.2.2 Reactivity Studies -
Huybrechts et al.^-26^ studied the oxygen inhibited photochlorination and
chlorine sensitized oxidation of perchloroethylene and pentachloroethane
at 353.3°K and 373.4°K. The reaction was initiated by light at 4358A in
which the rates of chlorination of both compounds were given. It was
observed that the rate of chlorination decreased progressively on adding
increasing amounts of oxygen and at the same time a chlorine photo-sensi-
tized oxidation took place. Analysis of products in experiments carried
to completion showed that 85% (±5%) of the oxidized perchloro and penta-
chloroethane appears as tichloroacetyl chloride and 15% (±5%) as phos-
gene corresponding to the overall reaction:
c2ci4 + Jgo2 > cci3coci
C2C14 + 0 2 > 2COC12
Trace quantities of carbon tetrachloride (.3%) and tetrachloroethylene
epoxide (.1%) were also detected. Horowitz et al.(27) studied the kinetics
of the photochemically initiated reaction between tetrachloroethylene
and pentane. Experiments were conducted at 25°C and energy was provided
by a high pressure mercury lamp. The reaction resulted in the formation
of pentyltrichloroethylene and HC1, and the rate constant for the rate-
determining step was estimated at 6.6xl02 sec"1. Goldfinger et al.*-2^
studied the gas phase photochlorination of perchloroethylene with
methane, methyl chloride, methylene chloride, chloroform, and penta-
chloroethane in a static system between 385° and 490°K, in which they
utilized light of different intensities in an effort to obtain "satis-
factory sensitivities." Reaction rates were followed monometrically
and photometrically allowing the following rate constant to be
measured:
C2C14 + Cl -->C2C15: log K = 9.4 where k is in units of
M"1 sec .
Uusoleil et al. 2^ investigated gas phase photochlorination of perchloro-
ethylene and pentachloroethane in a static system in which light of
11
-------
4047 - 4358A was utilized, and the reaction temperature was 418.80°K.
They also conducted other experiments at temperatures up to 564°K. There
is no documentation of smog reaction studies performed on perchloroethyl-
ene.
3.2.3 Ambient Measurements -
Available data on ambient background concentrations of perchloroethylene
are in conflict. Tohura KuriyangC^O) reported 1 ppm C2C14 as ambient
background concentration. His reported data was obtained using a spe-
cially designed non-dispersive ultra violet absorber. However, Reid et
al.C ) reported detection and measurements of ambient concentrations of
C2C14 below 1 ppm. Williams(3^ utilizing an electron capture detector,
reported background atmospheric concentrations of perchloroethylene to
be less than 1 ppb. Murrayet al.^1 ) made ambient measurements of £-2^4
in rural areas of Britain and away from large towns in which an attempt
was made to examine the environmental distribution of the aliphatic
chloro-compounds in which significant amounts of these compounds were
detected, particularly in air from towns and cities, as would be expected.
Simmonds et al. *• ^ conducted measurements of atmospheric halocarbons
throughout the Los Angeles area in which he reported average concentrations
of 0.125 ppb of perchloroethylene. Lillian and Singh^) have reported
background concentrations of 0.12 ppb C2C14 in the New Brunswick area.
This was determined by gas phase coulometry.
3.3 Methyl Iodide
5.5.1 Sources, Uses and Toxicity -
Methyl iodide is a colorless liquid with a sweet etheral odor. On
exposure to light a light brown color develops due to the liberation of
iodine. It is used mainly in the laboratory as a methylating agent and
as an intermediate in the chemical industry. It has been used in fire
extinguishers, as a vesicent in chemical warfare, in the process of
stone carving and it is used in microscopy as an embedding agent due to
its high index of refraction. Methyl iodide is prepared by the reaction
of dimethyl sulfate and sodium iodide in a still ^33 J. An old commercial
method involved the burning of seaweedP4,35)_
Methyl iodide is not used extensively and consequently, few cases of
poisoning have occurred. One was reported in 1901, in a dye factory^).
The symptoms were vertigo, diplopia, ataxia, and the patient swaying as
if drunk. When delirium and mania ensued, treatment in a mental hospital
was necessary but the mental dullness persisted. In 1940, an exposure in
a methyl iodide factory caused nausea, vomiting, diarrhea, vertigo, slur-
red speech, visual disturbances, tremor, drowsiness and then coma.
Another reason few poisonings have occurred is that methyl iodide is lac-
rimatory. The principal basis for methyl iodide's threshold limit value
of 5 ppm is its tendency to sensitize the skin.
12
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Methyl iodide is considered ten times as acutely toxic as carbon tetra-
chloride by the American Conference of Governmental Industrial
Hygienists(37) . Figure 1 includes the sales of CHjI for the period 1967
to 1973. Based on studies in the literature CIO, 16 j ancj on our own sampiing
(38) } the typical concentration of CH3I in air is 0.001 ppb.
From the sales figures and the scarcity of ambient data, one can only
estimate the residence time. It is obvious that methyl iodide has a
natural source. Lovelock et al.(10) measured methyl iodide in and over
the Atlantic Ocean. They suggest that the sea is the source of atmos-
pheric methyl iodide. The atmospheric residence time has been estimated
at 50 hours suggesting an annual production of 90 billion poundsClO). ^
possible reaction of atmospheric sea salt and methyl radicals can be the
source of CHsI. Biological methylations in seaweed and marine algae
should atso be considered as a possible natural source of methyl iodide.
3.3.2 Reactivity Studies -
Investigations of the photochemical reactions of methyl iodide are well
documented in the literature. Bates and Spence(39) were the first to
look into the photooxidation of gaseous methyl iodide. Their work, noting
that the absorption spectrum of methyl iodide (Figure 2) was continuous
above 2000A, indicates that decomposition is the primary effect of light
absorption. They varied the concentrations of oxygen from 10 to 600 torr,
methyl iodide from 10 to 100 torr, and light intensity from 30% to 100%
at 2537A". The results of their experiments point to a linear relation-
ship between light intensity and reaction rate.
2 A 1
dt (I2) k2 + (02) k3
where A = I0 (1 - e-a
I0 = intensity at 2537A
and a = molar extinction coefficient at 2537A
kj, k2 and kg being the rate constants determined from the reactions:
1) CHsI + hv ----- > CHs + I
2) CH3 + I2 ----- > CH3I + I
3) CH3 + 02 ----- > CH20 + OH
The Quantum yield was found to be 2.
At these relatively high concentrations, side reactions of the products
interfere with the elucidation of the quantum yield of the primary photo
lysis. Their experimental data fit the following expression for the
decomposition rate of methyl iodide:
[CH5I] = K(1.e-0.047[CH5I]
dt 12.3 + [02]
13
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40
o>
o
LU
_l
<
CO
Figure 1
U.S SALES OF HALOCARBONS
1967 - 1973
30
20
ro
O
co
LJ
10
to
1967
1973
-------
I
o
o
••*
m
o
o
"CN
ro
; O
r-o
O
: O
O
-co
CM O
c
•H
i
O
- O
tN
rn
I
a 6oq
15
-------
West and Schlessinger studied the photodecomposition of methyl
iodide vapor at 2537A. They found low yields of methane, ethane, ethylene,
and methylene diiodide. Their data changed the rate found in previous
study because Bates and Spence assumed no iodinated hydrocarbons as
products and simplified their reaction mechanism accordingly. The low
yields of products were accounted for by the great tendency towards
recombination of methyl radicals and iodine atoms. In the presence of
silver foil the yield of ethane increased forty times due to the
abstraction of iodine as silver iodide.
The major reactions of their mechanism were:
1)
2)
3)
4)
5)
6)
7)
CH3I +
CH3 +
CH3 +
CH2I2
CH2 +
CH3 +
hv >
I >
CH3I >
CH2 >
CH3 >
CH3
CH3I
CH4
CH2
CH2
C2H4
C2H6
+
+
12
+
I
CH2I
12
Iredale^ questioned the low yield of iodine and reacted methyl iodide
with nitric oxide. Nitric oxide was used to detect the presence of
radicals in the reactions of organic molecules. The compound CH3NO seems
to be removed from the reaction system as a solid and takes no part in
further chemical reactions at room temperature.
Absorption of light by methyl iodide to form an excited CH3I* molecule
was ruled out. His reaction scheme was:
CH3I + hv ----- > CH3 + I
CH3 + NO ----- > CH3NOi
CH3 + I2 ----- > CH3I + I
CH3 + I ----- > CH3I
Other reactions leading to methane and ethane were excluded due to the low
frequency of occurrence.
f 421
Leighton et al.v J photooxidized methyl iodide and studied the kinetics
of the product formation. They found discrepancies between the reaction
schemes proposed by earlier workers due to experimental errors. They
hypothesized the formation of a peroxide intermediate instead of an
hydroxide radical proposed by Bates and Spence '^9).
1) CH3I + hv ----- > CH3 + I
2) CH3 + I2 ..... > CH3I + I
3) CH3 + 02 ----- > CH302
4) CH302+ CH3 ----- > CH3OH + CH20. . .
They found the rate of methyl iodide decomposition:
16
-------
I[02]
dt ky
2k3 [I2] + [02]
Where the maximum quantum yield is unity.
This discrepancy in quantum yields was researched further by Iredale and
McCartney (-43) ^ They found that only when the nitric oxide pressure is
lower than 50 torr do the back reactions: CH3I + I2 > CH3I + I
CH3I + I > CH3I
begin to lower the quantum yield. When the pressure of nitric oxide is
greater than 50 torr the quantum yield approaches unity. Their experi-
ments, however, were performed over a spectrum from 2500 to 2700A, and
weighted averages had to be employed in their intensity calculations.
Shultze and Taylor' J postulated the "hot" methyl radical formation in
the approximately 50 - 60 kcal/mole greater than the carbon-iodine bond
strength of 54 kcal/mole. The reaction mechanism they proposed was:
1) CH3I --kV_CH3* + I
where CH3* represents a methyl radical with excess energy of 32 kcal/mole.
2) CH3* + CH3I > CH4 + CH2I
3) CH3* + M > CH3 + M
4) CH3 + CH3 > C2H6
5) I + I > I2
6) CH3 + I2 > CH3I + I
Reaction (2) is the sole source of methane production. They found an
apparent activation energy for methane production of zero in the tempera-
ture range of 40° - 100°C. Ordinarily 8 kcal/mole of energy would be
needed to produce methane. Hence, an increase in temperature would vary
the amount of methane produced. This is not the case. The quantum yield
of methaae production is a low 0.003 due to the fact that reaction (3) takes
place more readily. Unless the reactants at the instant of collision are
oriented in the right way, the "hot" radical will merely lose kinetic
energy in collision with a third body (3). The methyl radical, having
lost its excess energy will most likely react in (6), but if iodine
can be removed sufficiently rapidly by fixation with silver as silver
iodide, reaction (4) would become appreciable.
r45")
Willard^ explains that a "hot" radical is one which has been born with
energy much in excess of the average energy of the surrounding molecules
and has not yet reached thermal equilibrium with them. If it reacts
chemically before becoming thermalized, the reaction is a "hot" reaction.
A typical hot radical can carry its excess kinetic energy through only a
few collisions.
Although hot radicals are probably produced in many photochemical dis-
sociations, only a few examples of their reactions have been reported.
This is due to the low quantum yields of hot reactions.
17
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The quantum yield for the reaction of hot methyl radicals with methyl iodide
is 0.003 for radicals formed at 2537A, whereas it is 0.03 for radicals formed
at more energetic 1849A light. Most photochemical reactions have been studied
under conditions where the rate of thermal reaction is sufficient to mask
the rate of a hot reaction. This means that a high light intensity source
is needed to produce observable hot radical reactions.
Martin and Noyes*- -* investigated the reactions of methyl radicals with
oxygen. They found formaldehyde to be one of the principal products of
methyl iodide and oxygen mixtures at room temperature. Ethane was not
found among the products when oxygen was present and quantities of carbon
dioxide and carbon monoxide result ultimately from reaction of methyl
radicals with oxygen.
Hudson, Williams and HamillA ) studied the moderation of hot methyl
radicals in the photolysis of methyl iodide. Moderators are inert gases
added to the reactions to physically decrease the rate of production of
a product, in this case, methane. The investigators used neon, argon,
helium and carbon dioxide as moderators and found methane production de-
creased by a constant factor, yet this decrease was independent of tem-
perature change. This experiment strengthened the hot radical hypothesis
of Shultz and Taylor.
Souffie, Williams and Hamill^ ' continued the study of hot radical reac-
tions. They measured rates of production of ethane, methane and iodine as
functions of pressure, temperature, iodine concentration, various moderators
and reactive substances added to obtain a better picture of the hot radical
reaction.
Their reaction mechanism:
1) CH3I + hv ---> CH3* + I
2) CH3* + CH3I --->CH4 + CH2I
3) CH3* + RH --->CH4 + R
4) CH3* + CH3I ---> C2H6 + I
5) CH3* + CH3I ---> CH3I + CH3
6) CH3* + M ---> CH3 + M
7) CH3 + CH3 —-> C2H6
8) CH3 + I2 ---> CH3I + I
9) CH3 + I —-> CH3I
10) CH2I + I2 ---> CH2I2 + I
11) I + I + M ---> I2 + M
The rates of production of ethane and iodine decrease during the early
stages of the photolysis. Addition of iodine further decreases the ethane
rate to a limiting rate. This limiting rate is independent of further
changes in iodine concentration. The authors attribute this production
of ethane to the hot radical reaction (4). The quantum yield of "hot
ethane) is 3.6 x 10'4.
The authors also suggest that "hot" ethane production is the result of
18
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excess translational energy in the methyl radical while methane production
may be the result of excess internal energy. In all their experiments,
the extent of decomposition of methyl iodide sample (present at 150-200
torr) was less than 5%.
Christie's work^SJ was to investigate the photochemical oxidation of methyl
iodide under conditions such that oxygen molecules were competing with io-
dine molecules and/or iodine atoms for reaction with methyl radicals, and
thus to obtain information about the kinetics of the reactions involved.
She found that in excess oxygen, the photochemical reaction was zero
order, i.e. the amount of iodine produced was directly proportional to the
time of irradiation. The quantum yield of iodine is 0.5 since her experi-
ments showed that oxygenated products did not contain iodine. The reaction
mechanism was:
1) CH3I + hv —-> CH3 + I
2) CH3 + I2 ---> CH3I + I
3) CH3 + 02 ---> products
4) CH3 + 02 + M —-> CH302 + M
5) I + I + M -—> 12 + M
6) CH3 + I ---> CH3I
where the reaction (6) is negligible compared with reaction (2) in the
methyl iodide plus oxygen system. Note also the reaction (4) is third
order.
McKellar and Norrish^J combusted methyl iodide under slow and explosive
conditions. Hydroxyl radicals and formaldehyde were among the products
formed. Their data indicates the reaction CH3 + 02 —> H2CO + OH occurs.
The significance of this reaction in the ambient atmosphere is slight,
however, since explosive conditions are highly unlikely.
Heicklen and Johnston( ^B reacted low concentrations (260 ppm) of methyl
iodide in oxygen with ultraviolet radiation above 2200A. Their measure-
ments were made by leaking the reaction mixture directly into the electron
beam of a mass spectrometer during photolysis. Their very complex reaction
mechanism involves 24 steps to include such reaction products as: 12, H2CO,
CH3OH, CH3OOH, H20, C02, CHOOH, CH3OOCH3 and CH3OI. A number of free
radicals are also involved in the system. The best way to present these
is in the following chart which shows the source and fate of free radicals
and source of products in the proposed complex mechanisnu *J.
19
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CH,I
COj.HCOOH
Photolysis of CH^I in §2 - Proposed Mechanism(51J
Doepker and Ausloos*- -* photolyzed methyl iodide at 2537A in the presence
of added organic materials. The "hot" methyl radicals reacted with these
additives to form methane. By keeping the ratio of additive to methyl
iodide sufficiently large (10:1), reactions of methyl radicals with methyl
iodide could become negligible so that the measured quantum yield of
methane be attributed to the reaction of a "hot" methyl radical with the
additive. These additives included CDsCHs, C2D2> C6H12> CH3OCH3, and
(CH3)CCD3. They found isotope effects to be of minor importance. They
also found that the quantum yields of methane decrease by a factor of 10
when the wavelength of the light increases to 3130A. No new reaction
mechanisms were proposed.
Mains and LewisC53) flash photolyzed methyl iodide at 2537A. They found
acetylene and ethylene among the products in much larger quantities than
had been observed in low-intensity photolysis. Since the acetylene yield
is strongly dependent on the iodine concentration, they proposed that a
two-stage process is involved: CH2l + ^2 —> ^2^~^2
CH2I + CH2I-I2 —> C2H2 + 2HI + I2
The methyl iodide radical was detected but they cite wall effects to adsorb
the hydrogen iodide which they could not find. They also suggest that the
"hot" methyl radical has an energy of 36 kcal/mole. 9.5 kcal/mole of this
energy is vibrational, 23 kcal/mole is translational, and the remaining
energy goes to the iodine atom. The translational energy of the methyl
radical is only slightly above the activation energy for the hydrogen
20
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abstraction reaction, CH3* + CH3I --- > CH4 + CH2l. This fact plus the
more efficient reaction, CH3* + M --- > CH3 + M account for the low methane
yields.
Gilbert, Lambert and Linett1 J thoroughly explored the photolysis of
013! + NO mixtures. Their mass spectrometer was sampling directly from
the reacting system. Their research emphasis was on the reaction pro-
ducts after the formation of CH3NO. The most noteworthy products were
HCN, OH, and HN03. Their reaction mechanism is shown below:
1) CH3I + hv ---> CH3 + I
2} I + I M — > I2 + M
3) CH3 + NO + M ---> CH3NO + M
4) 2CH3NO ---> (CH3NO)2
5) CH3NO + NO 5 = i CH3CNO)2
6) CH3(NO)2 + NO ---> CH3 + N2 + N03
7) N03 + NO ---> 2N02
8) N02 + CH3 + M ---> CH3N02 + M
9) N02 + CH3NO ---> CH3N02 + NO
CH3NO ---> CH2NOH
11) CH2NOH ---> HCN + H20
12) CH2NOH + NO + hV ---> HCN + HN02 + H
13) H + HX ---> H2 + X
14) H + NO + M ---> HNO + M
The fact that hydrogen cyanide and nitric acid have been identified as
reaction products of CH3I + NO mixtures is not significant when one
realizes that the total pressure in the reaction chamber did not exceed
three torr. These conditions do not exist in the environment.
In summary, the continuous abgorption of gaseous methyl iodide in the
wavelength region around 2500A indicated that the primary quantum yield
of the dissociation CH3I + hv — > CH3 + I was unity, but the quantum
yields of the decomposition products, mainly methane, ethane, and iodine
are very small. This low quantum yield is attributed to the back reactions
reforming methyl iodide. Increasing sensitivity of analytical techniques
has enabled workers to show that the photooxidation of methyl iodide is
complicated and a large number of reactions occur.
The forward reaction is accelerated by the presence of silver which
reacts with atomic or molecular iodine and by the presence of oxygen and
nitric oxide which react with methyl radicals. A striking feature of the
photodecomposition of methyl iodide was the discovery that the reaction
CH3 + CH3I --- > CH4 + CH2I would require an activation energy of 8-9 kcal/
mole for it to make an appreciable contribution to the overall reaction at
room temperature. The "hot" methyl radical theory was propagatgd to ex-
plain these results. The energy of the absorbed photon at 2537A is greatly
in excess of the H3C - I bond dissociation energy.
21
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5.3.5 Ambient Measurements -
CHjI has been measured in the ambient air by Lovelock as mentioned
earlier, at a mean background level of 0.001 ppb. However, some serious
question concerning experiment artefacts in the analytical technique
have been raised by Wilkness et al.( •* • The implication of this possibility
to some of our measurements will be commented upon in a later section
(6.2.2) of this report.
5.4 Carbon Tetrachloride
5.4.1 Sources, Uses, Toxicity-
Carbon tetrachloride is an excellent solvent. It is volatile and, when
present in high concentrations, has a narcotic effect on the central ner-
vous system^-*). Carbon tetrachloride is a clear, colorless, sweet
smelling liquid which had been used as an anesthetic, anthelmintic, and
as a hair shampoo 0>°). These uses have been discontinued and the greatest
use of carbon tetrachloride today is in industrial applications: metal
degreasing; solvent for rubber cement and paints; fumigant for grain;
dry cleaning agent; manufacturing intermediate for soap, chloroform,
dyes, insecticides, plastics, inks, and freon gases; and as a fire extin-
guishing agentf56). In the 1930's, carbon tetrachloride displaced
naphtha as the premier dry cleaning agent due to its non-flammability.
It was soon displaced by tetrachloroethylene for its lower toxicity.
One of the major uses of carbon tetrachloride has been as a fire extin-
guisher. It appeared with the advent of electrical machinery since it
is non-conductive. Carbon tetrachloride fire extinguishers were found
to inhibit fires of flammable liquids. Injuries had been reported where
workers were overcome by carbon tetrachloride vapors and a by-product
produced when carbon tetrachloride is exposed to high temperatures, phos-
gene gas.
Carbon tetrachloride is an extremely toxic substance, capable of causing
extensive damage to the liver, kidneys, lungs, and the heart. The three
modes of exposure are ingestion, inhalation, and absorption through the
skin and mucous membranes. The most frequent route of exposure appears
to be by inhalation. Carbon tetrachloride poisoning causes necrosis in
both the liver and kidney. The kidney damage results in destruction of
cells whose function, in part, is to eliminate toxins from the blood. A
major function of the liver is also to detoxify the blood. When carbon
tetrachloride poisoning is severe enough to cause kidney and liver damage,
the body's exposure to this poison is thus prolonged, which compounds its
toxic effect to the rest of the body.
As an example of how small an amount of carbon tetrachloride can produce
dangerous concentrations, one teaspoonful of carbon tetrachloride placed
in an unventilated bathroom 6' by 10' by 8? would vaporize and produce a
concentration of 100 ppm. The threshold limit value of carbon tetra-
chloride set by the American Conference of Governmental Industrial
22
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Hygienists is 5 ppm. It is for this reason that the Federal Government
has banned its use in the home as a hazardous substance (^7) _
The total sum of carbon tetrachloride in the atmosphere greatly exceeds
that which is accounted for by global production figures . It was there-
fore postulated by Lovelock et al.™' and later confirmed by Singh et al .
(• J that a natural source exists for carbon tetrachloride in the atmos-
phere. Lovelock^ ' states that carbon tetrachloride could be formed by
reaction of chlorine with methane in the troposphere. Carbon tetrachloride
has been experimentally observed by us( ) during the simulated tropo-
spheric irradiations of standard mixtures of perchloroethylene in air as
vill be discussed later in this report (Section 7.2). It is even pos-
sible that tropospheric reactions of other chlorinated olefins would
result in the formation of carbon tetrachloride in the troposphere.
5.4.2 Reactivity Studies -
The reactions of carbon tetrachloride have been thoroughly investigated
in the literature. Its photochemical reactions, however, are less well
understood.
In 19,35, Lyons et al . ( ) photo-oxidized liquid carbon tetrachloride at
2537A. They proposed the reaction:
2CCl4(aq) + 02 -by- COCl2(g) + 2Cl2.(g)
This reaction only occurs with the exclusion of water and it does not
react in the absence of oxygen. The appearance of phosgene is significant
in that phosgene has been used as a nerve gas and as a vesicant in
chemical warfare.
In 1961, Ung and Schiff^59^ reacted oxygen atoms with carbon tetrachloride
vapor. Their primary reactions were bimolecular:
CC14 + 0 ---> COC12
---> CO + 2C12
No other products were reported except for carbon dioxide which could be
produced by the chain sequence:
CO + Cl ; — - COC1
COC1 + 0 — -> C02 + Cl
COC1 + 02 — > C02 + CIO
The rate constant for the primary process was found to be independent of
atomic oxygen, molecular oxygen and carbon tetrachloride concentration:
K = 3.3 x 10~14 ex ~4° cm3 molecule"1sec"1.
exp
This could lead to a carbon tetrachloride half-life of 4600 years at the
ambient levels of atomic oxygen in the troposphere.
23
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Carbon tetrachloride has been used as an intermediate for the manufacture
of many industrial chemicals. Hexachloroethane is formed when CC14 and
alcohol are reacted in ultraviolet light ^ '. The byproducts of this
reaction are hydrochloric acid, chloroform, and various aldehydes.
Styrene and vinyl butyl ether are photopolyjnerized in the presence of
carbon tetrachloride (I) at 3660A" and 3080A. These processes occur
under controlled conditions of temperature and pressure that are not
present in the troposphere.
Carbon £etrachloride was photolyzed and reacted with ethane and ethylene
at 2537A by Roquitte and Wijnen^ '. the primary step produces trichloro
methyl radicals and chlorine atoms. Although only 1% of the carbon
tetrachloride originally present was decomposed, this reaction could
have0serious consequences in the stratosphere where radiation below
2900A is not filtered by the ozone layer ( ) .
The following reaction mechanism was suggested to explain the formation
and recombination of the products from the photo-initiated reaction
with ethane:
1)
2)
3)
4)
5)
6)
7)
8)
CH14 + hv
Cl + C2H4
2C2H4C1 -
2C2H4C1 -
2C2H4C1 -
2CC13 —
C2H4C1 +
C2H4C1 +
> CC13 + Cl
---> C2H4C1
-> (C2H4C1)2
--> C2H4 + C2H4 + C2H4C12
--> C2H5C1 + C2H3C1
> C2C16
CC13 > CC13C2H4C1
CC13 ---> CC13H + C2H3C1
The reaction mechanism for the reaction with ehtylene is very similar to
1-8 and the rate constants are usually faster due to the lower activation
energies of the ethyl radical.
Tomkinson et al.C ) decomposed carbon tetrachloride at 2537A but not at
3130A. This study of the abstraction of chlorine atoms by methyl radicals
was made at high temperature (90 - 140°F) . Since the source of methyl
radicals was a thermal decomposition of di-t-butylperoxide, errors can
be introduced in the determination of the reaction rates and especially
in trying to extrapolate the data to ambient conditions.
Clark and Tedder ( J reacted hydrogen atoms with carbon tetrachloride and
generated the products: hydrogen chloride, chloroform, methylene dichlor-
ide and methyl chloride. The experimental apparatus did not include a
device for measuring hydrogen atoms from the discharge of a Wood's tube,
but the author's estimate the concentrations of the reactants to discover
the reaction mechanism. The hydrogen atoms produced in the troposphere
(from the photolysis of formaldehyde) react rapidly with oxygen to form
hydroperoxyl radicals C '0.
H + 02 + M — > HOO + M + 60 kcal/mole
24
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It is unlikely then, that the hydrogen abstraction reaction:
CC14 + H ---> CC13 + HC1
would be significant in the environment.
In 1968, Lautenberger et al.t J photochemical ly reacted n-butylamine with
liquid carbon tetrachloride. The products were mainly chloroform, hexa-
chloroethane and ammonia. Approximately 90% of the light was at 2537A and
the reaction cell was maintained at room temperature.
Both ionic and free radical mechanisms have been proposed to explain their
experimental results but the conclusion of the authors is that a free radi-
cal chain reaction is initiated by the absorption of light by the charge-
transfer complex of the amine with carbon tetrachloride:
RRNH2 + CC14 + hV ---> RNH2 + Cl- + CClj.
In the presence of oxygen, chloroform and hexachloroethane yields are
reduced. The trichloromethyl radicals (CC13') react with oxygen to yield
phosgene which then reacts quickly with excess amine to form urea.
CC13- + 1/2 02 ---> COC12 + Cl-
CC13- + 02 -— > COC12 + Cl-0
COC12 + 2RNH2 ---> CO(NHR)2 + 2HC1.
These photochemical reactions in which carbon tetrachloride vapor is de-
composed preclude reactions in which carbon tetrachloride could participate
in the ambient atmosphere. Experiments were conducted to ascertain the
possibility of any tropospheric sinks for carbon tetrachloride with special
emphasis on nitrogen oxides, free radical hydrocarbons, ozone and atomic
oxygen .
3.4.5 Ambient Measurements -
Carbon tetrachloride has been found injjeasurable amounts wherever and
whenever it has been sought. Lovelock^ reported the mean aeriaL con-
centration over the Atlantic Ocean at 0.071 ppb. Simmonds et al.f J
reported CC14 in the range 0.1-1.63 ppb. during a three-day study in
17 ppb.
Lovelock
has recently reported "background" concentrations of 0.11 and 0.14
ppb. in Western Ireland and over the North Atlantic respectively. He
also gave data for CC14 concentrations as a function of altitude over
central England which showed values of about 0.07 ppb. at 10 km. - i.e.
above the tropopause.
4 .-. . -
the Los Angeles Basin. ,We have reported urban levels of about 0.
in New Brunswick, N.J.( J and 0.19 ppb. in Bayonne, N.J.f- J. L
( ^ has recentl reorted "backround" concentrations of 0.11 and
5.5 1:1:1 Trichloroethylene
5.5.1 Sources, Uses and Toxicity -
Methyl chloroform has become an increasingly popular solvent in recent
25
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years due to its low toxicity. It has been used primarily for metal de-
greasing but it has been promoted as a carbon tetrachloride substitute
for use as an intermediate.
1,1,1-trichoroethane is a colorless liquid possessing a chloroform-like
odor. It is not flammable and has been used in hairsprays, dry cleaning
and as an aerosol propellant in combination with freons. The principal
toxic action is the depression of the central nervous system. Below the
current threshold limit value of 500 ppm, no physiological effects have
been observed. Above that concentration, 1,1,1-trichloroethane has been
used as an anesthetic. This is not a unique property of methyl chloroform,
but one common to most chlorinated aliphatic hydrocarbon solvents.
1,1,1-trichloroethane has also been used as a degreening agent for produce
such as tomatoes, berries, applies and citrus fruits. This appliation
produces a deep, even coloration characteristic of normal ripe fruit and
it effectively retards decay due to mold and fungus growth( .-) . For this
specific agricultural application, methyl chloroform has been exempted
from the requirement of a tolerance level for residues(68).
Experiments with humans exposed to 1,1,1-trichloroethane vapor show that
most of the compound is exhaled immediately. 1,1,1-trichloroethane vapor,
in high concentrations (>900 ppm), acts as an anesthetic agent and has
only a slight capacity to cause reversible injury to the liver or kidneys
5.5.2 Reactivity Studies -
There is a dearth of literature concerning the photochemical reactivity of
1,1,1-trichloroethane. It does not absorb in the visible or ultraviolet
range. The C-C bond dissociation energy for 1,1,1-trichloroethane is
84.2 kcal/mole based on the heats of formation of chlorinated methyl
radicals^70). This amount of energy is much too high to achieve any lysis
of the bonds in the troposphere.
A study of chlorine-photosensitized oxidations of chloroethanes at 4360A
was conducted in 197l(71). A concentration of 60 torr of 1,1,1-trichloro-
ethane was irradiated with chlorine and oxygen at 355°K no products
were formed. The authors explained this on the basis of the extreme exo-
thermicity of the dissociation of the hypothetical radicals involved.
This compound seems to be tropospherically inert. The low solubility of
1,1,1-trichloroethane suggests that the ocean is not a sink. It may re-
main in the ecosphere as residue or more likely flux into the stratosphere.
5.5.5 Ambient Measurements -
Like 0014, CHjCC^ is considered to be ubiquitous and is generally found
at sub-ppb levels. Simmonds et al. ^-^ reported Cti^CCl^ levels of 0.01 to
2.3 ppb during their 1974 study in the Los Angeles area. Lillian and co-
workers' values for New Brunswick(•H) and Bayonne, New Jersey^ ) were
26
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respectively of 0.27 and 0.83 ppb. Lovelock's^ -* "background" concentrations
were 0.065 and 0.075 ppb in western Ireland and over the Atlantic.
5.6 Trichloroethylene
5.6.1 Uses. Production and Toxicity -
Trichloroethylene is widely used in industry as a solvent particularly in
the vapor degreasing of fabricated metals where 95% of that produced is
utilized. In addition, TCE is used as an extraction solvent, a chemical
intermediate and an anesthetic.
A common and simple technique for the production of trichloroethylene is
through the use of acetylene as the raw material, but in recent years
ethylene has been used instead due to its lower cost. For the most part,
two methods are employed for the production of TCE from ethylene. One is
photochlorination of 1,2 dichloroethane followed by dehydrochlorination of
tetrachloroethane. The second technique involves oxychlorination of
ethylene or dichloroethylene at high temperatures.
The demand for the production of TCE has been kept active mainly due to
its use in the aircraft and space industries as well as other metal
industries. The market was hurt in 1966 by a report by the Los Angeles
County Air Pollution Control District which claimed that TCE was a sig-
nificant pollutant contributing to the formation of photochemical smog
( '^). Since then, several private studies(21Jwere undertaken which
tended to downgrade its role in smog formation and restored some confidence
in its use. The future of TCE in industry is as of now highly uncertain.
The graph of sales of TCE (Figure 1) compiled from U.S. Tariff Commission
reports on Synthetic Organic Chemicals tends to substantiate the uncer-
tainty of TCE's future role in industry. Since 1970 the production of
TCE has dropped some 29%.
The major physiological responses which result from exposure to trichloro-
ethylene is one of central nervous system depression. Symptoms of acute
exposure may include nausea, mental confusion and fatigue. Several
fatalities have occurred due to severe acute exposure. These have been
attributed to ventricular fibrillation resulting in cardiac failure.
Also possible effects to liver metabolism are suspected. In a recent
study by the National Cancer Institute*- '$ ', trichloroethylene has been
found to cause cancer in mice. The evidence submitted is being questioned
by producers of TCE due to the large doses given the mice and the means
by which it was given Congestion) which isn't the principal mode by which
it is transferred to humans. Due to this evidence, an alert has been
issued.
The toxic hazard ratings of trichloroethylene compiled by Sax*- ' is as
follows:
Acute local - Irritant - slight
Ingestion - slight
27
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Inhalation - slight
Acute systematic - Ingestion - moderate
Inhalation - high
Skin absorption - moderate
Chronic local - Irritant - slight
Chronic systematic - Inhalation - slight
Skin absorption - slight
Due to the above hazard rating and other information not mentioned, a
TLV of 100 ppm has been set for trichloroethylene in the industrial
environment .
5.6.2 Reactivity Studies -
There seems to be no doubt that trichloroethylene is a photochemical ly
reactive compound. Several investigators have conducted studies on the
compound utilizing different conditions yielding different results.
Several of these studies will be discussed in this section.
Huybrechts and Meyers -* studied the chlorine photosensitized oxidation
of the compound at 363° and 403°K. They found that, depending on the
temperature, between 82 and 90% of the compound was oxidized to dichloro-
acetyl chloride. Other products observed to form included chloroform,
phosgene, carbon tetrachloride, carbon monoxide, HC1 and trichloroethylene
epoxide. The mechanism they proposed to explain this is as follows:
»• 2C1
Cl + C2HC13 + C2HCl4
C2HC14 + C12 •* C2HC13 + Cl
C2HC14 + C2HC14 -»• chain termination
C2HCl4 + 02 -> C2HC1402
C2HCl402 •> C2HC14 + 02
C2HC1402 + C2HC14 ->• C2HC1402C2HC14
C2HC1402 + C2HC1402 -> C2HC1402C2HC14 + 02
C2HC1402 + C2HC1402 -»• 2C2HC140 + 0~
C2HC140 -> C2HC130 + Cl
COC12 + Cl
2
CC13-CHC10 -*• CC13 + HC1 + CO
4-Cl2
CC14 + Cl
CO + HC1 + CO
V&2
CHC12-CHC10 •* CHC12 + COC12
IC12
CHC13 + Cl
28
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In addition, Huybrechts and Meyers'- ' studied the oxygen effect in the
photochlorination and found that oxygen tended to inhibit the reaction.
They determined that the greater the 02/C1 ratio the smaller the extent of
photochlorination and the favored reaction is the oxidation.
Mayo and Honda^ J studied the autooxidation of trichloroethylene in the
liquid phase and obtained results similar to Huybrechts, with the dif-
ferences being accounted for by the fact that they didn't work in the gas
phase.
Smog studies concerning trichloroethylene were conducted by Katherine
Wilson( ; after the Los Angeles Air Pollution Control District first
implicated TCE as a significant contributor to photochemical smog and
producers became alarmed. In the study she found that TCE was not as
significant as the Los Angeles team suspected. She concluded that if
the ambient concentration didn't climb significantly upwards, no problem
would result from TCE as far as smog is concerned. The conclusion was
supported by experiments in which higher concentrations than those found
in ambient air were used. The photoreactivity was found to be negligible.
In a smog forming atmosphere, the presence of TCE didn't significantly
increase eye irritation; and under conditions similar to those when
pollutants leave an industrial stack (100 ppm), no increase in oxidant
level was detectable. ,
( 78 ~>
Bufalini et al. ^ -1 conducted experiments which showed phosgene as a
product as well as traces of formyl chloride. Other products noted to
form included HC1, carbon monoxide, HCOOH and nitric acid. The initial
reactants in this irradiation were trichloroethylene and NC^.
3.6.3 Ambient Measurements -
Ambient measurements of TCE made by various investigators have documented
its presence in both air and surface waters. Murray and Riley^ J recorded
concentrations of TCE in Great Britain and over the northeast Atlantic
Ocean. The concentrations recorded varied significantly from one site
to the next. At rural stations the values varied from 2-28 ng/m^; in
the city of Liverpool the value reported was 850 ng/m^; over the north-
east Atlantic detected concentrations varied from 1-22 ng/m^; and the
surface waters analyzed contained from 5 - 11 ng TCE per liter. Lovelock
(j monitored for trichloroethylene over the north Atlantic and western
Ireland. The mean concentrations reported were 15 parts per 10,1^ volume
and less than 5 parts per 10-^, respectively. Lillian et al.C ) reported
maximum and minimum values of TCE in the northeast United States from
Baltimore, Maryland to Whiteface Mountain, New York. The mean values
recorded at the different sites varied from highs of .92 ppb and .71 ppb
in Bayonne, New Jersey and New York City, New York, respectively; and
lows of less than .05 and .1 ppb in Baltimore, Maryland and Whiteface
Mountain, New York, respectively.
29
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3.7 Ethvlene Dichloride
5.7.1 Uses, Production and Toxicity -
Approximately 75% of all the ethylene dichloride produced in the U.S. goes
into the production of the vinyl chloride monomer; 7% is used as a chlori-
nated solvent intermediate such as in the production of trichloroethylene,
and 5% is used as a lead scavenger. Other than the above, ethylene di-
chloride is used as an industrial solvent in cleaning and extraction
processes, as a fumigant and as an anti-knock additive to gasoline.
Ethylene dichloride is produced by the addition of ethylene and chlorine.
The techniques which are used in industry to carry out this reaction are
as follows: (1) A reaction of ethylene and chlorine in a body of ethy-
lene dichloride in a stirred tank reactor, or (2) the contact of ethylene
and chlorine gases with a circulation stream of ethylene dichloride in a
packed or empty reactor.
The demand for production of ethylene dichloride in industry has increased
sharply since the mid 1960's as can be seen from the plot of sales compiled
from U.S. Tariff Commission Reports (Figure 1). This increase is largely
attributed to its role in the production of the vinyl chloride monomer
and its increased use in the manufacturing of chlorinated solvents. As a
solvent it has been displaced by trichloroethylene and and CC14. Due to
increasing concern over the vinyl chloride monomer's role as a carcinogen,
one cannot accurately predict future trends.
Physiological responses which result from exposure to high concentrations
,of ethylene dichloride are effects on the eyes, upper respiratory tract
and the nose. In addition, liver, kidney and adrenal injuries may occur.
Symptoms of acute exposure include irritation to the eyes, nose and upper
respiratory tract, gastrointestinal upset, mental confusion, nausea and
dizziness.
The toxic hazard ratings compiled by Sax^ ' is as follows:
Acute local Irritant - high (may cause permanent damage)
Ingestion - high
Inhalation - high
Acute systematic - Ingestion - high
Inhalation - high
Chronic local - Irritant - moderate (reversible and irreversible
changes, not severe)
Chronic systematic -Ingestion - high
Inhalation - high
Skin absorption - moderate
At present the TLV for ethylene dichloride is set at 100 ppm but due to
30
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the information available at the present time the A.C.G.I.H, has recom-
mended a lowering of the TLV to 50 ppmC^7 ).
3.7.2 Reactivity Studies -
Ethylene dichloride does not absorb wavelengths above 2900A, No instances
of its photochemical reactivity in air in this region of the spectrum have
been found in the literature.
5.7.3 Ambient measurements -
As far as is known, no ambient measurements of ethylene dichloride have
been reported,
3.8 Vinyl Chloride
3.8.1 Sources, Uses and Toxicity -
Vinyl chloride (monochloroethylene) is made by oxychlorination of ethylene
or by hydrochlorination of acetylene. It boils at -13.9°C, and at room
temperature is a colorless, faintly sweet-smelling gas. As a gas it is
flammable and explosive, but is usually handled as a liquid under pressure.
It has been used as a propellant for aerosol products.
Vinyl chloride is the parent compound of a series of thermoplastic resin
polymers and copolymers which find wide use in containers, wrapping such
as for food packing materials, film, electrical insulation, water pipes,
conduit, and a variety of other industrial and consumer products. In the
United States vinyl chloride has been made commercially since 1939. The
vinyl chloride industry is divided into three segments: monomer production,
polymer production, and fabrication. Vinyl chloride (VC) is used primarily
in the production of polyvinyl chloride (PVC), a resin which is producer
through batch processing.
The vinyj. chloride/polyvinyl chloride industry is a multi-billion dollar
concern involving several dozen large manufacturers of vinyl chloride and
polyvinyl chloride resin and thousands of fabricators producing a variety
of products based on polyvinyl chloride and probably employing more than
300,000 workers. The production of vinyl chloride from 1969 to 1974 rose
at an average annual growth rate of 14%. In 1975 output dropped from a
1974 level of 5.60 billion Ib. to 4.20 billion lb., a 25.4% drop in pro-
duction. An increase in output above 1975 levels is predicted for 1976.
Because of the renewed increased consumption of polyvinyl chloride, the
lack of new production capacity and the effect of regulations on vinyl
chloride emissions, there may be a production shortage by 1977.
Prior to January 1974, vinyl chloride had been regarded as having moderate
liver toxicity. Patty^ 7™ stated that the dominant handling problem
was the fire and explosion hazard. On January 22, 1974, the Occupational
Safety and Health Administration (OSHA) was informed by the National
Institute for Occupational Safety and Health (NIOSH) that the B.F. Good-
rich Chemical Company, the largest U.S. producer of PVC resin, reported
31
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that deaths of several of its employees from a rare liver cancer (antio-
sarcoma) may have been occupationally related, A joint inspection of the
B.F. Goodrich plant by State and Federal agencies followed soon after,
and on February 15, 1974, a fact-finding hearing was held.
The investigating groups (in the United States and Europe) looking into
the circumstances of VC and PVC production began one of the most urgent
medical inquiries in history. In the United States they included: (1)
the Occupational Safety and Health Administration (OSHA) of the Depart-
ment of Labor; (2) the National Institute for Occupational Safety and
Health (NIOSH); (3) the Environmental Protection Agency (EPA); (4) the
Center for Disease Control (CDC); (5) the National Institute of Health
(NIH); (6) the American Cancer Society (ACS); (7) the Environmental
Science Laboratory, Mount Sinai School of Medicine, New York; plus some
industrial investigations, in particular those undertaken by the
Manufacturing Chemists Association (MCA).
In 1975, vinyl chloride production in the United States was 4.20 billion
Ibs. The estimated loss of vinyl chloride during processing is
approxiately 3.0 - 6.3% and the polyvinyl chloride loss is of the order
of 1.3% at the PVC polymerization facilities., with more than 75% of the
losses being vinyl chloride air emissions^°°i Thus, in the United States
alone, substantial amounts of vinyl chloride (in 1975 probably exceeding
120 million Ibs.) and large quantities of polyvinyl chloride (probably
exceeding 50 million Ibs.) were discharged into the environment during
the PVC polymerization production processes. Most of the vinyl chloride
escapes directly into the atmosphere, with lesser amounts dissolved in
water effluent streams and entrapped in sludge and solid wastes. The
polyvinyl chloride losses occur as particulates in air emissions, sus-
pended solids in water effluents, and components of solid wastes.
The principal source of vinyl chloride leakage is in the operation of
the polymerization kettles. Other losses occur during the transfer of
vinyl chloride from tank cars to storage areas, during the PVC drying
process and from a variety of valves, flanges and pump seals. Dust
collector losses and disposal of oversize particles are some of the
areas accounting for polymer losses.
Another area of potential vinyl chloride exposure concerns the presence
of the unreacted monomer entrapped in polyvinyl chloride products. The
existence of residual vinyl chloride in articles made from vinyl chloride
polymers is related to the manufacturing process and the physical struc-
ture of the polymers. In early March 1975, Professor Cesare Maltoni
disclosed that vinyl chloride caused cancer when fed to rats (81. ) . These
results put vinyl chloride in conflict with the "Delaney" clause. This
states that any substance that causes cancer in even one animal must be
banned from contact with food or cosmetic products.
On September 3, 1975, the U.S.F,D.A. proposed regulations to restrict the
uses of vinyl chloride polymers in contact with food*- 82) by banning
rigid and semirigid polyvinyl chloride articles intended to contact food,
32
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including bottles, blister packs, boxes and pipe. The use of PVC in
coatings applied to fresh citrus fruits would also be prohibited^ )t
The continued use of PVC in food contact items such as films, gaskets,
capliners, flexible tubing, and package coatings would be permitted by
FDA. The use of water pipes made from PVC would also be allowed, sub-
ject to provisions concerning an interim food additive regulation.
Until recently, levels of unreacted vinyl chloride entrapped in PVC
products were estimated between 50 and 1000 ppm^O ) . The new methods
of making PVC are aimed primarily at reducing worker exposure to vinyl
chloride as required by OSHA. According to the Society of the Plastics
Industry, the residual level of vinyl chloride in certain resins is now
less than 1 ppm(84 )^ an(j due to improved manufacturing methods, for all
practical purposes, no vinyl chloride can migrate to foodC ).
The production of PVC bottles accounts for only 8% of the total plastic
bottle market. Anticipation of federal regulations has led some PVC
makers to shift away from PVC bottles for food products(°° ). Packagers
have been concerned about the safety of PVC bottles since the FDA dis-
covered in 1973 that vinyl chloride leaked into distilled alcoholic
beverages contained in PVC bottles(82 ).
Vinyl chloride has been used as an aerosol propellant in a variety of
products, including hair sprays. Although the precise data when vinyl
chloride was first used for this purpose is uncertain, as far back as
1959, a review on the use of CH2CHC1 and CH2C12 as propellants in aero-
sols was publishedC ). At one time, vinyl chloride was one of the most
commonly used propellants in aerosol products C ). A list of over
50 different pesticidal sprays containing vinyl chloride including many
used indoors has been compiled by EPA(°^' ) .
A 1964 report estimated that an aerosol product sprayed in a tiny room of
282.5 ft3 for 30 sec. would result in 0.025% (or 250 ppm) vinyl chloride
by volume(88 )_ it might be noted here that Dr. Cesare Maltoni reported,
in early 1975, on the development of angiosarcoma of the liver in test
animals at vinyl chloride levels as low as 50 ppmC J.
In April 1974, FDA asked for recalls of hairsprays containing vinyl
chloride made by two major producers. Other manufacturers were requested
to recall any outstanding stocks of such products from the market. EPA
had earlier asked makers of 23 insecticide sprays containing vinyl
chloride to reformulate their products ( 90 ). The use of vinyl chloride
as a propellant in cosmetic aerosols was prohibited in a final regulation
C 91 ) t A new drug application for the marketing of aerosol drugs con-
taining vinyl chloride as a propellant was also required.
The EPA banned the use of vinyl chloride as a propellant in certain
pesticide aerosols in April 1974(^2) anc[ the Consumer Product Safety
Commission banned the use of other self-pressurized household products
containing vinyl chloride in August 1974^ J,
33
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As results of the various toxicity studies of vinyl chloride became known,
the mounting evidence pointed to this chemical as the main agent respon-
sible for the induction of liver angiosarcoma. Much of the early toxicity
studies were conducted by Professor Cesare Maltoni, Institute Di Oncologia,
Bologna, Italy, and by Dr, P,L, Viola of the University of Rome. In 1970
almost four years before the discovery at the B.F. Goodrich plant, Dr.
Viola produced tumors in the bones, skin and lungs of rats exposed to
30,000 ppm (3%) of vinyl chloride( '. In 1971, Dr. Maltoni began to
test animals at lower concentrations of the gas.
His investigations involved various types and levels of exposure to vinyl
chloride. At the New York Academy of Sciences meeting, May 10-11, 1974,
he reported the development of angiosarcoma of the liver and other types
of tumors at levels of atmospheric exposure as low as 50 ppm^ ).
Vinyl chloride has also been linked to other health-related effects. Con-
siderable data exists concerning the toxic effects of vinyl chloride from
atmospheric exposures, especially by inhalation and occupational contact.
There is ample evidence that vinyl chloride inhalation can result in acute
toxicity manifested by an array of symptoms including unconsciousness.
Vinyl chloride has been implicated as a causative agent for acroosteolysis
of the hands and feet as well as systemic effects among workers engaged
in the manufacture of vinyl chloride. In one report, approximately 1-3%
of the workers involved in the manual cleaning of PVC reactor vessels had
at one time experienced acroosteolysis( ). It is characterized by a
soreness and thickening of skin at fingertips, gradual dissolution of
bone calcium at the fingertips and toes, skin sores, and frequently
heightened sensitivity of hands to cold. These symptoms occur only after
several years of high levels of exposure.
A number of chemical companies, in early 1974, just after the B.F. Goodrich
disclosure, stated that they had not been in violation of the federal
standard for exposure to vinyl chloride vapors. The Treshhold Limit Value
or TLV had been set at 500 ppm, in April 1959, by the A.C.G.I.H. Dow
Chemical, a major vinyl chlo,rj.de maker, as early as 1961, had urged use
of a 50 ppm exposure levelC ). Studies in 1961 on the chronic effects
of vinyl chloride on laboratory animals(96) indicated some effects at
100-200 ppm and 100 ppm was suggested as a T.L.V.
Partly because of toxicity studies of vinyl chloride implicating it as the
causative agent in the induction of angiosarcoma of the liver, and because
employees were being exposed at levels around the experimentally observed
effect level of 250 ppm, an Emergency Temporary Standard (ETS) was pro-
mulgated by OSHA on April 5, 1974( 97) . This standard reduced the level
from a ceiling of 500 ppm to 50 ppm.
Additional vinyl chloride toxicity studies by the Industrial Bio-Test
Laboratories, Northbrook, Illinois revealed that mice exposed to 50 ppm
vinyl chloride for 7 months developed angiosarcoma of the liver. On
May 10, 1974 OSHA proposed a permanent standard for employee exposure at
34
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no detectable level, as determined by a sampling and analytical method
capable of detecting vinyl chloride at concentrations of 1 ppm with an
accuracy of 1 ppm ± 50%. Engineering controls and work practice
programs were also required to reduce levels below detectability. On
October 4, 1974 the new federal standard was published in the Federal
Register and became effective on April 1, 1975f- J. The standard
required that exposure to vinyl chloride not exceed 1 ppm averaged over
an 8-hour period and 5 ppm averaged over a 15-minute period.
5.8.2 Reactivity Studies -
Because of the interest in occupational exposures to vinyl chloride,
there have been several recent studies on various aspects of its
photooxidation, Sanhueza et al.(" ) have studied the chlorine-atom
and 0(3?)-initiated oxidations of this compound. Gay et al.C ) have
also investigated vinyl chloride as part of an overall study of chlori-
nated ethylenes. Here, vinyl chloride was U.V.-photooxidized in air in
the presence of N02. Ozonolysis was also studied to help elucidate the
contribution of this process to the reaction mechanism. Huie et al.
(100) investigated the kinetics of the reaction of 0(3p) with several
haloethylenes including vinyl chloride. Arnold et a.l.( *) have
examined chemiluminescence from the reaction of oxygen atoms and vinyl
chloride. RennertC*"^ studied vinyl chloride photolysis in presence
of radical scavengers (12, HI, H2S) in presence of added inert gases.
Williamson and CvetanovicC ^ have measured reaction rates of ozone
with chlorinated olefins in CC14 solution. Ozone attack was found to
be electrophylic, decreasing in rate as the number of chlorine atoms in
the molecule increased.
All these studies show that vinyl chloride is reactive, though patterns
of products and postulated reaction mechanisms vary, depending on the
particular experimental conditions employed. Arnold et al.f J obser-
ved formaldehyde and many other products, and suggest a mechanism
involving 0-atom addition to the double band, forming an intermediate
which decomposes to an aldehyde and a carbene. Williamson and
Cvetanovic(103) observed only phosgene as a product. Huie et al
failed to detect formaldehyde as a product of 03 reacting with vinyl
chloride. They attributed this to reaction of the formaldehyde with
03 to produce formic acid or CO and H20.
Gay et al.C ^8) have detected formaldehyde in other products. For-
maldehyde was produced rapidly, rising to 0.85 ppm after 160 minutes.
An ozonolysis study of vinyl chloride produced a maximum of 0.75 ppm
formaldehyde. They also observed formation of HC1 at levels about 50%
higher than the formaldehyde. Formyl chloride was also a product, but
being thermally unstable it rapidly decomposes to CO and HC1. Sanheuza
and HeicklenC ") in a study of the 0(^P)-initiated oxidation of vinyl
chloride, also found evidence of formyl chloride production in the
presence of oxygen.
35
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Arnold et al.^ ' observed vibrationally excited HC1 when vinyl chloride
is introduced into a stream of 0-atoms. This was attributed partly to
thermal decay of formyl chloride, but predominantly by a reaction of
hydrogen atoms with Cl2 (another observed reaction product),
3.8.3 Ambient Measurements -
The EPA and other groups have made ambient measurements of vinyl chloride.
Monitoring results from EPA at a number of industrial PVC and VC plants
seem to indicate that the levels of vinyl chloride in ambient air near the
plants fluctuate sharplyr^Nearly all collected samples contained detectable
levels of vinyl chloride, <1 ppm. A number of individual air samples con-
tained >1 ppm. Water effluents were also monitored for vinyl chloride, and
the levels varied considerably. The highest level was 20 ppm, with 2-3 ppm
a more typical level. Grimsrud and Rasmussen( ^ ) searched for vinyl
chloride plus 18 other halocarbons in the atmosphere of the rural north-
west (southeast Washington state). Vinyl chloride could not be detected,
<5 ppt, using a GC-MS system.
5.9 Dichloromethane
3.9.1 Sources, Uses and Toxicity -
Of the more than 520 million pounds of dichloromethane produced during
1973, most of this was applied in uses requiring a very strong solvent.
Dichloromethane, in addition to those uses already cited, is also used
as a solvent for cellulose esters, for polyvinyl acetate, chloride and
chloroacetate, methyl and ethyl cellulose, rubber, bitumen, pitch, oils,
waxes and other resins^ J. It is especially suitable as a solvent in
the preparation of polycarbonates by phosgenation in the presence of
pyridinef ). Dichloromethane is also finding increasing use instead
of trichloroethylene in vaDor degreasing of metal parts in industrial
metal fabricating plants( J. This action stems from the recently
reported liver carcinogenicity of TCE in mice. An additional use now
just being realized for dichloromethane involves the use of CH2C12 instead
of previously used TCE in the manufacture of some ground and instant
decaffeinated coffees^ ). According to the FDA, methylene chloride
use in these processes is allowed since it continues on the approved list
pending further studies. FDA has no information at present to indicate
that CH2C12 poses any hazard.
Another important use of methylene chloride is as an aerosol propellant.
Some reviews covering its manufacture, physical, physiological and chemi-
cal properties, and its use in aerosols have been publishedC '10"J.
Cannizaro et al.C °) resolved 18 volatile components of commercial
aerosol formulations, consisting of mixtures of hydrocarbons, chlorinated
and fluorinated hydrocarbons, and aliphatic halides. The properties and
methods of preparation of CH2C12 in aerosol formulations has also been
published by K. Bergwein(m) .
36
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The suitability of dichloromethane as an aerosol propellant is dependent
upon a number of factors. These are: (1) vapor pressure between 15 -
100 psig at 70°F; (2) non-flammability; (3) freedom from odor; (4)
chemical inertness to other materials in the formulation; (5) a low
order of toxicity and (6) freedom from irritation1- 12> . While it is a good
solvent, dichloromethane is also slightly irritating to the eyes when
sprayed in the air. As a propellant, dichloromethane can be found in
formulations such as shaving cream, dry shampoos, deodorants, hair cos-
metics, pharmaceutical products, food aerosols and pesticides.
Other health effects of dichloromethane have been observed. Its vapor
offers a definite industry hazard unless care is taken with ventilation.
It is narcotic and toxic though not as much as chloroform which it closely
resembles in its physiological effects. At high concentrations it is a
rapid anesthetic and death may result from asphyxia. Its high vapor
pressure at ordinary temperatures increases this hazard. Extreme care
with ventilation and avoidance of skin contact is necessary where this
solvent is used industrially. Long or repeated exposure to lower sub-
narcotic concentrations may result in organic injury to liver, kidneys
and other organs <- J .
The in-home use of paint removers containing dichloromethane results in
the absorption of this solvent, which is metabolized to CO, and can
eventually lead to heart stressCH'*). jn Sweden, human subjects were
exposed for controlled periods to varying concentrations of dichloro-
methane, but no detrimental effects were observed. Methyl ene chloride
has been an air pollutant problem in other areas also, ranging from air
contamination of space vehicles to a hazard during welding operations
*• ** •
3.9.2 Reactivity Studies -
Some simulated atmospheric photodecomposition studies have been reported.
Dilling et al.C116) found that less than 5% of the initial dichloro-
methane had reacted in 21 hours in the presence of NO, and less than 5%
in 7.5 hours in presence of N02- Sanhueza and HeicklenC^S ) studied the
chlorine-atom sensitized oxidation of dichloromethane and chloromethane
at 3655A and 32°. The initial major products were found to be CHC10 and
HC1, with smaller amounts of CC120 and CO being produced. The formation
of CO resulted from the decomposition of CHC10, thus it was considered a
secondary product.
Arnold et al.' investigated the infrared chemiluminescence from the
reactions of oxygen atoms with halomethanes. In these reactions, the
chemiluminescence produced arose from vibrationally excited CO and HC1 .
Some other reaction products observed were C02, C12, CC120, C2H2C12,
C2H2C120, C2H2 and Qty. It seems likely that the initial reaction steps
are the abstraction of a hydrogen atom from the halomethane
0 + CH^^ ---> OH + CH3_nXn (x=Cl or Br)
37
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followed rapidly by
0 + OH = H + 02 + 16.8 kcal/mole
Lin(118) has observed in chemical laser studies that 0("lD) atoms obtained
from the U.V, photolysis of ozone readily insert into the C-H bonds of
molecules such as CHnX4n (X=F, Cl) and that the resulting unstable alcohol
eliminates HX (X=F, Cl) producing vibrationally excited HF or HC1.
5.9.5 Ambient Measurements -
Some groups have attempted to measure CH2C12 in air and water. Grimsrud
and RasmussenC ) using a Durapak Carbowax 400/Porasil F column linked
to an MS system, failed to detect CH2C12 in the northwest United States,
CH2C12 <5 ppt. The authors concluded that the detection of CH2Cl2 and
CHsCl might increase further the relative role chlorocarbons play as the
means of transporting chlorine to the stratosphere. The compound
however was detected at a level of about 530+30 ppt. Dowty et al.(
detected CH2C12 and many other halogenated aliphatic compounds in water
from a New Orleans water treatment plant.
5.10 Polychlorinated Biphenyls (PCBs)
3.10.1 Sources, Uses and Toxicity -
Polychlorinated biphenyls (PCBs) are one member of a group of industrial
organochlorine compounds which have recently been recognized as a wide-
spread, persistent contaminant in the environment. Since their intro-
duction in 1929, PCBs have come into use as dielectric fluid in capacitors,
insulating/cooling fluids in transformers, and high temperature and
pressure lubricating fluids'12u). However, it was not until 1966 that
Jensen first identified PCBs in environmental samples(121)_ The usually
present "unknown interfering peaks" on gas chromatograins of DDT extracts
were identified as PCBs. Following Jensen's discovery, many surveys of
fish and birds were conducted revealing the ubiquity of this new contam-
inant (122). The potential hazard PCBs presented to human health was
emphasized in 1968 by the sickening of more than 1000 persons in Japan
who had used a cooking oil contaminated by PCBs. Again, in 1971, the
U.S. Food and Drug Administration seized 75,000 eggs laid by chickens
which had eaten PCB contaminated fish meal. In reaction to the opinion
of the scientific community, Monsanto Chemical Company, the sole producer
of PCBs in the U.S., voluntarily announced a restricted sales policy for
PCBs(123). Beginning in 1971 Monsanto would sell only to closed system
users and Monsanto would recycle or incinerate expended PCB containing
material. Nevertheless PCBs are still being measured in air, water,
sewage, and in human adipose tissue.
As evidenced by its structure, substitution of chlorine atoms onto the
ten available sites of the biphenyl ring can produce many isomers.
38
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Two hundred and ten combinations are possible but about 100 are thought
to be probable considering reaction conditions and the unlikelihood of
greatly unbalanced chlorination of the benzene rings(124).
Sissons and Welti used a combination of gas chromatography, NMR, and mass
spectroscopy to examine the isomeric content of several commercial PCB
mixtures(125)_ Identification of nearly all of the mixtures was accom-
plished. They found that Aroclors (Monsant's trademark for PCBs) of
42% chlorine content contained 45 isomers, those of 54% chlorine con-
tent contained 69 isomers and those of 60% content contained 78 isomers.
The appearance of the PCBs differ according to their chlorine content.
Aroclors of 21 to 62% chlorine appear as colorless mobile oil to light
yellow sticky resin(^°). Aroclors of 68 and 70% chlorine are white and
white crystalline powders respectively. The PCBs as a group and
especially the higher chlorine content mixtures are extremely nonflam-
mable. Coupled with their special dielectric properties, the PCBs have
become invaluable in electrical application. The higher chlorinated
compounds possess a very low vapor pressure. Most of the PCBs are
soluble in organic solvents but not in water. Solubilities in water as
stated by Monsanto are as follows: 42% chlorination - 200 ppb, 48%
chlorination - 100 ppb, 51% chlorination - 50 ppb, and an estimated 25
ppb for 60% chlorination. In addition, the PCBs are extremely resistant
to chemical attack. They are unaffected by moderate to rigorous attack
by acids or bases; but PCBs will react with sodium hydroxide at elevated
temperatures to form phenolic compounds. Lastly, Monsanto states that
PCBs are resistant to mildew and microorganisms.
The chemical inertness of PCBs combined with their low volatility, flame
resistance, high dielectric constant and their compatibility with chlor-
inated hydrocarbons find PCBs many specialized applications. Before
Monsan-to adopted its "closed system" use sales policy in September 1970,
60% of sales were for dielectric material in electrical capacitors and
for insulating/cooling oils in large power transmission transformers
(^ ), Another 25% of sales were for use as plasticizers, which
included the incorporation of PCBs in polyvinylchloride for flexibility
and also for us« in producing carbonless reproducing paper, e.g. credit
card receipts. Less than 5% of sales went for miscellaneous applications
including; surface coatings, adhesives, printing ink and pesticide
extenders.
As previously mentioned, Monsanto initiated a sales policy whereby they
would phase out sales to customers of open system use and would continue
to sell to users where the end product would be recycled or destroyed.
39
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This amounted to reducing their sales for largely electrical capacitors
and power transformers. (In addition Monsanto set up a system to col-
lect fluids from closed system uses and to either regenerate or destroy
PCBs through high temperature incineration(*^),) By 1972 Monsanto
expected to have reduced their sales to 100% closed system use.
Lastly Monsanto would move to discontinue the higher chlorinated bi-
phenyls and move to biphenyls of 48% or less chlorine content. The lower
chlorinated biphenyls would stand a better chance of microbial degradation
in the environment as supported by some experimental results.
In addition to Monsanto, other companies outside the U.S. produce PCBs.
Worldwide production figures are not available. However, the U.S. pro-
duced 33 x 106Kg in 1970 while Japan produced an estimated 11 x 106KgC )
Japan's market was similar to the U.S.: 40-50% for capacitors, 15% for
transformer oil, 10-15% for heat transfer fluid, 5% for plasticizors,
15% for carbonless reproducing paper, and 5-10% for export.
Foreign manufacturers of PCBs and their trade names include: Prodelec,
France (Phenocolor); I.G. Farber Industrie A.G., Germany (Clophen);
Kanegafuchi Chemical Company, Japan (Kannechlor); Mitsubishi - Monsanto
Japan CSan-To therm); Caffaro, Italy (Fenclor); and U.S.S.R. (Sovol)
Of the estimated 45 x lO^Kg of PCBs produced in North America between
1930-1970 it has been reasoned that 78% was released to the environment;
the balance of the material is either still in use or has been destroyed
by incineration (1™) The major releases are estimated to be SxlO^Kg to
the air,6xlo7Kg to fresh and coastal waters and 27xlO?Kg to dumps and
landfills. The chlorinated biphenyls are believed to be concentrated in
several major areas: (1) buried in landfills and dumps; (2) attached to
sediments in rivers and lakes, (3) attached to sediments on the continen-
tal shelf and (4) distributed over land and sea by aerial fallout and
disposal from ships.
5.10.2 Reactivity Studies -
One possible route of degradation for PCBs is through photochemical reac-
tions . Pesticides have been shown to undergo photochemical transformation
to sometimes even more toxic substances than their precursorsf ' ).
Two halocarbon pesticides, aldrin and dieldrin, have been shown to undergo
photo-isomerization via both sunlight and laboratory irradiationsC ^.
The products were 2-3 times more toxic than the original compounds.
Examination of the ultraviolet absorption spectra of a few chlorinated
biphenyl isomers shows that the onset of absorption occurs around 310
nanometers^ '. (See Figures 3 and 4). The absorption of the bi-
phenyl structure beginning in the longer wavelengths and continuing to
its first major absorption peak between 240 nanometers to 260 nanometers
(depending on the chlorine substitution) is generally attributed to
40
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FIGURE 3
ULTRAVIOLET ABSORPTION SPECTRA OF SO'ffi CHLORINATED BIPHEJYLS.
o
i
1
III
IV.
V. 2,1,6,2'
200
220
1 -—-CACKLOPOBlFrr/.'YL
o - -^ 28o
300
...TTH IIAIJ
41
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GUjRi; 4
ULTFAVIOLLT ABSORP'JIC i SrlCTRA OF TWO MCNOCHLOFOBIPHSJIYLS,
I, U-CKLOECBIPHIirfL: RE,
II.
200 220
2liO 260 280
WAVEIEIIGTH
42
300 320
-------
excitation of the IT electrons of the phenyl ring(l*°). The appearance of
this TT IT* transition peak is clearly defined in the non-ortho substituted
isomers but becomes less distinct with increasing ortho-substitution. The
isomers of a given series also shows that the ortho substituted isomer
exhibits a molar absorptivity of an order of magnitude less than its cor-
responding meta and para compounds. Figure 4 shows little difference
between 3,3-dichlorobiphenyl and 4,4'-dichlorobiphenyl (Note: absorption
peak at 257 nanometers) while the 2,2'-dichlorobiphenyl lies at most a
full order of magnitude lower. This effect is believed due to the
inability of the biphenyl structure to achieve planarity between the
rings in the ortho substituted compounds. The relationship of this
phenomenon to the excited states and to the reactivity will be discussed
later.
Most photochemical studies of PCBs have been conducted in the liquid phase.
Usually irradiations have been carried in solvents such as benzene, or
light aliphatic alcohols. Concentrations were relatively high (100-1000
ppm). The quantity and purity of UV radiation was generally unspecified
except to note that wavelengths were greater than 250 nanometers.
Compounds irradiated include the para substituted chloro, bromo and iodo
biphenyls in benzene, chlorobenzene in ethanol and in isoproponol, and
0-halonapthalenes in benzenef ). The results of these studies
support a general mechanism for the photolysis of aromatic halides. The
absorption of a photon produces an excited species which either returns
unreacted to its ground state or homolytically cleaves the carbon-halide
bond. This results in the production of a phenyl radical and a halogen
atom.
Ph-X + hv ---> Ph-X* ---> Ph- + X •
Ph = phenyl, biphenyl or phenyl
napthalene
X = I, Br, Cl, F (para substitution)
The direct reaction of the Ph-X* species with another reactant is not
believed to occur. The products observed are due to the resulting reac-
tion of the aromatic radical and the halogen atom; hence the products
observed are dependent on the nature of the solvent and its impurities.
Reactions of the aromatic radical observed include polymerization,
hydrogen abstraction, and radical scavenging by oxygen.
The products observed upon irradiation of 4-iodobiphenyl in benzene serve
to illustrate the pathways open to the photodissociation products. The
identified products include: biphenyl, p-terphenyl, iodine, and 4-hydro-
zybiphenyl( » '. The distribution of these products was dependent
on the degree to which oxygen was initially present in the solvent. Upon
removal of oxygen, the hydrozybinhenyl decreased along with an increase
in the biphenyl/terphenyl ratio(142), The observed effects are explained
by the following mechanisms:
43
-------
1) C6H5-C6H4I + hv >C6H5C6H4< • + I"
2) C6H6 + C6H5C6H4- -—>C6H5 - C6H4HX^> (A)
3) C6H5 - C6H4+ (A) —->C6H5C6H5 + C6H5C6H4C6H5
4) (A) + \ 12 —->C6H5C6H4C6H5 + HI
5) C6H5C6H4 + HI —->C6H5C6H5- + I
6) 2 I- >I2
7) (A) + 02 -—>C6H5C6H6C6H5 + H02 •
8) C6H5C6H4- + 02 —->C6H5C6H400-
9) C6H5 - C6H400- + HI-->C6H5C6H4OOH --->C6H5C6H4OH
Similar results axe obtained upon in
or isopropanol( ^. These solvents
poundsC ). Thus, upon cleavage of
Similar results 1ax£ obtained upon irradiation of chlorobenzene in ethanol
are excellent hydrogen donating com-
of the carbon-chlorine bond, the phenyl
radical can easily abstract an alpha hydrogen from the alcohol. This re-
sults in the quantitative formation of benzene and some hydrogen chloride.
Small amounts of ortho, para, and meta chlorobiphenyl were formed also
when ethanol was the solvent. No phenols were observed since the sol-
vents were deoxygenated or saturated with high purity nitrogen.
The foregoing limited knowledge of the photochemistry of chlorinated bi-
phenyls coupled with the current concern over PCBs in the enviornment has
prompted a few investigators to conduct studies specifically aimed at
elucidating the photochemistry of PCBs in the environment. These studies
published since August 1971 through June 1974 are mainly qualitative in
nature; they were conducted to determine mechanisms and products. All
but the most recent do not attempt to determine product balances or
kinetics. The studied PCBs include both individual pure isomers (symet-
trically substituted di, tetra, hexa, and octa isomers) and a 54% chlori-
nated biphenyl commercial mixture. The sample varied in concentration from
1 ppm in solution to some solid film studies. Two investigations involved
vapor phase irradiations. The sample conditions for the most part were
not typical of environmental conditions; the solvents were chosen mainly
for their ability to solubilize the isomer and for their ultraviolet
transparency. Some of the studies used sunlight but most used mercury
vapor lamps having a maximum ultraviolet energy output at about 310
nanometers which is normally found in the troposphere. However, most of
the studies do not define the intensity at these wavelengths, hence, com-
parison among experimenters or normal tropospheric conditions is impossible.
Irradiation times vary from 5 minutes in some laboratory studies to 2
months under natural sunlight.
The conditions and results of the various studies are compiled in Table 3.
44
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In spite of the qualitative nature of these investigations, some genera-
lizations can be drawn concerning the mechanisms operative in the photo-
chemistry of PCBs. Irradiations of individual isomers generally yielded
mechanistically interpretable products while irradiations of the complex
commercial mixtures could only be roughly characterized.
Although as previously noted, rates and percent disappearance of the
isomers are not comparable between experimenters due to lack of quanti-
fication of light fluxes; experiments conducted with a common light
source are useful. When this was the case, rates of disappearance of PCB
mixtures were found to differ according to the solvent used(147). irradia-
tion of 2 ppm of 54% chlorinated biphenyl mixtures were found to disappear
at decreasing rates in the order: hexane, water, and benzene^1*^. This
order is believed to be related to the ability of the biphenyl radical to
abstract a hydrogen from the solvent. This order is an anomaly since water
has the highest H bond energy of any of the three. The insolubility of
PCBs in water may contribute to an apparent removal.
The number of chlorines and the position of the chlorines on the biphenyl
structure are expected to have an effect on the rate of disappearance. It
is expected that due to steric effect, the ortho substituted compounds
should eliminate the ortho chlorine faster than the corresponding non-
ortho compound. However, upon irradiation at 310 nanometers for 24 hours,
the following isomers in hexane (1000 ppm) were still present at the
indicated percentages: 3, 4, 3', 4'-tetrachlorobiphenyl - 32%, 2, 6, 2',
6'-tetrachlorobiphenyl - 29%, 2, 5, 2', 5'-tetrachlorobiphenyl - 33%,
2, 4, 5, 2', 4', 5'-hexachlorobiphenyl - 3.8%, and 2, 3, 4, 5, 2', 3',
4', 5'-octachlorobiphenyl - 1%(**°). Thus, the only distinguishable
effects are that the higher chlorinated compounds tend to degrade more
efficiently. The effect of the position on the ring is not apparent.
The most frequently observed result of irradiation of PCB mixtures or
individual isomers is stepwise dechlorination. The removal of chlorines
is seen in different solvents with both sunlight and 310 nanometer mer-
cury vapor light. Dechlorination, for example, occurred upon irradiation
of 2, 4, 6, 2', 4', 6'-hexachlorobiphenyl in hexane where penta, tetra,
trik di and monochlorobiphenyls in addition to biphenyl were formed(^ ).
Dechlorination is observed along with isomerization and chlorination in
situations where one or less hydrogens are available for abstraction from
the solvent and no other species are present to react with the dissociated
components of the biphenyl. This situation occurs when PCBs are irradi-
ated as pure solids or in degassed solvents such as a chlorofluorocarbon.
In this environment the phenyl radical and the chlorine atom have no
choice but to recombine or react with other PCB molecules. Thus dechlori-
nation and chlorination occur leading to isomerization.
Under prolonged irradiation (90-100 hours) polymers are formed even at
concentrations as low as 10 ppm^^^J. Since chlorination and dechlori-
nation products can again participate in the same dissociation processes
48
-------
the biphenyl radical can be regenerated,
1) ClvPh Ph C1 + hv+ Clv Ph Ph C1Y , + Cl' (dissociation)
A. A. •"• A."" JL
2) Clx Ph Ph Clx_]_ + Cl- -»- Clx Ph Ph Clx (recombination)
3) Clx Ph Ph Clx + Cl(+H)cix Ph Ph Clx+1 (chlorination)
4) Clx Ph Ph-Cl.^} 4 A •> Clx Ph Ph C1X-1A (Radical Fraction)
A = Reactant e.g. H, 02, NO, €'tc.
Thus the products of these reactions may still be able to further under-
go photochemical reactions. But polymerization reactions may also be
competing.
5) 2 Clx Ph Ph-Cl^} •»• Clx Ph Ph (Clx_j) (Clx)Ph Rh (Cl^)
In most of the irradiations where oxygen was not removed from the solvent
oxygenated products were observed. These were roughly characterized by
thin layer chromatographic techniques as being hydroxy and carboxy
compounds. When one group of investigators irradiated in methanol they
detected 80 compounds via gc/ms( ). Most of the products were oxygen-
ated chloroaromatics - among them a series of methoxy substituted
chlorodibenzofurans with 1-5 chlorines and a series of p-methoxy PCBs.
Recent studies have centered upon the nature cf the excited electronic
states, the mechanisms and the rates of different PCB isomers of the
same chlorine content. Nordblum and Miller irradiated 4,4'=dichloro-
biphenyl at 310 nanometers in a variety of organic solvents(150). When
the 4,4'-dichlorobiphenyl was irradiated in both degassed methanol and
2-propanol, 4-chlorobiphenyl and hydrogen chloride were formed in
quantitative yield. The new hydrogen on the 4-chlorobiphenyl was abstia* -
ted from the methyl of the methanol as determined by using deuterated
methanol. This implies a typical free radical reaction where the weaker
C-H bond is broken in preference to the OH bond. Consistent with similar
studies it was found that the rate of reactant disappearance was slowed
by the presence of oxygen and the product mixture was complex.
The second of the recent studies by Ruzo, Zabik, and Schuetz is a
systematic study of the rates and medianisms of the photo recution of
6 tetrachloro biphenyl isomers^ ' '. Again, irradiations at 310
nanometers in both hexane and methanol produced hydrogen chloride,
dechlorination products accounting for 90 - 95% of the reacted starting
material and between 1 and 3% methoxy substituted biphenyls in the case
of the methanolic solutions. The PCBs studied are summarized below.
PCB Designation
3,4,3',4' - tetrachlorobiphenyl I
2,4,2',4' - tetrachlorobiphenyl II
3,4,3',4' - tetrachlorobiphenyl III
49
-------
2,3,2',3' - tetrachlorobiphenyl IV
2,5,2',5' - tetrachlorobiphenyl V
2,6,2',6' - tetrachlorobiphenyl VI
The products formed after 20 hours irradiation in hexane and methanol
are tabulated below:
PCB Dechlorinated Product Methoxylated Products
I 3,4,4' - trichlorobiphenyl trichloromethoxy biphenyl
4,4' - dichlorobiphenyl
II 2,4,4' - trichlorobiphenyl trichloromethroxybiphenyl
4,4' - dichlorobiphenyl dichloromethroxybiphenyl
4 - chlorobiphenyl
III 3,5,3' - trichlorobiphenyl
IV 2,3,3' - trichlorobiphenyl trichloromethoxybiphenyl
2,3,2' - trichlorobiphenyl dichlorodimethoxybiphenyl
3,3' - dichlorobiphenyl
V 2,5,3' - trichlorobiphenyl trichloromethoxybiphenyl
3,3' - dichlorobiphenyl dichlorodimethoxybiphenyl
3 - chlorobiphenyl
VI 2,6,2' - trichlorobiphenyl trichloromethoxybiphenyl
2,2' - dichlorobiphenyl
Examination of the products formed, clearly indicates that ortho chlorines
yield products resulting from the loss of these. In their absence, meta
chlorines were eliminated and para were the last to cleave. The methanol
formed products by attacking at the leaving chlorine position.
Quantum yields of reactant disappearance were also determined for these
compounds. These also showed a strong correlation with the position of
the chlorines. Compounds I and III, which are non-ortho-substituted,
exhibited the least efficiency for reaction. This is believed to relate
to the preferred planar geometry for the excited triplet^ '; thus the
sterically hindering ortho chlorine leaves preferentially to relieve the
strain.
Ruzo et al, also found the triplet to be the excited stateC ). In
cyclohexane, the efficiency of intramolecular energy transfer was 100%
(0 isc = 1.0 0,05] for compounds I, II and III, The initial absorption
leads to an excited singlet which then quickly leads to the first trip-
let,
1) [A] + hv > [A]*
2} X[A]* > 3[A]*
Since the quantum yields are the same for both ortho and non-ortho sub-
stituted compounds the process is apparently not affected by planarity
of the biphenyl,
50
-------
The intersystem crossing yield of 1.0 indicates that there are no com-
peting processes for the singlet, i.e., fluorescence, internal conver-
sion. The resulting triplet state is believed to be the reactive species
leading to either deactivation, dissociation, or direct reaction. The
lifetimes of the triplets also show a correlation with the position of
the substituent chlorines. The non-ortho substituted compounds here
demonstrate a lifetime 3 times longer than the corresponding ortho sub-
stituted compounds; but they also demonstrate a rate about 10 times
slower to react than the ortho substituted compounds and a deactivation
rate constant 3 times slower than the ortho compounds.
Ruzo et al. proposed the following scheme for the reactions they observed.
27) Cl
28)
29)
30) Cl
3D
32)
Cl +HC1
+C1
Cl
51
-------
Presumably the reaction of the biphenyl radical with oxygen would also
account for the observed phenolic compounds found by other investigators;
it should be noted, however, that if collision of the excited species
(the triplet) with oxygen occurs before dissociation of the oxygen an
excellent triplet quencher will lead to the deactivation of the triplet
to the ground state.
5.10.5 Ambient Measurements -
A recent review of published measurements of PCBs in surface waters found
PCBs quite ubiquitous. A background level for unpolluted fresh waters
averaged about 0.5 ppt.f J; the Great Lakes averaged about 5 ppt.
Moderately polluted rivers and bays yielded a common value of 50 ppt.,
e.g. the Milwaukee River. Very polluted Japanese and U.S. rivers
averaged about 500 ppt.C153).
Few direct measurements of PCBs in air have been made. EPA measured
between 1 to 50 nanograms/M^ in 4 major U.S. citiesQ^l) % Another sur-
vey conducted by Woods Hole Oceanographic Institute found an exponential
decrease in concentration of PCBs with increasing distance from the
Boston-New Jersey industrial complex(154)_ Concentrations ranged from a
high of 5 Ng/M-5 at Vineyard Sound to 0.05 Ng/M^ two thousand kilometers
out over the north Atlantic. These values reflect the sum of vapor and
particulate samples and are uncorrected for collection efficiency. Other
studies have indicated the presence of PCBs in air indirectly. A study
of organochlorine pesticides in rainwater in the British Isles found
PCBs present in all samples collected from 7 stations over a 12 month
period^ ). Concentrations were not determined. A similar study in
the south of Sweden noted PCBs present in aerial fallout ranging from
600 - 10,000 Ng/M3/month(156).
Extensive surveys of PCBs in fish and birds demonstrate a clear tendency
for accumulation and magnification in food chains. The highest concen-
trations of PCBs are found in the fish eating and scavenging birds
(herons, terns, eagles) and in marine bird predators (falcons). The
highest concentrations in marine systems are found in the fat of the
top predators such as porpoises, sharks, and seals (6 - 1800 ppm).
Laboratory experiments have shown that aquatic invertebrates and fish
can accumulate 3 x 10^ and 7 x 10^ higher than ambient water; the top
predators may be as much as 10' higher than ambient water. If fish-
eating birds are involved then magnification as high as 10^ or 10^ is
likelyC121).
Surveys of the occurrence of PCBs in man have been conducted by EPA. Of
2189 samples analyzed PCBs were found at 1 - 2 ppm in 30% of the adipose
tissue samples( ^, Mean levels in human milk were of the order of 60
ppb; blood plasma averaged about 2 ppb,
Research into the toxicity of PCBs has been complicated by the variation
in composition from manufacturer to manufacturer and even from batch to
52
-------
batch; in addition some commercial mixtures are known to contain toxic
contaminants such as p-chlorodioxins and chlorinated dibenzofurans(121).
Low level feeding of PCBs to rats demonstrates a tendency for accumulation
in the fat ^ , Once feeding has ceased as long as 240 days may be
required until the concentration in the fat reaches a peak. A half-life
of 30 days has been observed in decreasing the concentration. Hydroxy-
lated metabolites have been observed. LDso for rats were of the order
of 5 - lOg/kg. No deaths have ever been directly attributed to PCBs in
man.
Chronic intake of small daily doses of lOmg/kg over 50 days causes
chloracne in man^^M. Reproductive effects have been observed in
chickens at 8 ppm continuous feeding and in fish at 0.9 ppb in ambient
water.
Release of PCBs to the environment from 1929 until 1970 has been largely
unregulated. The American Conference of Governmental Industrial Hygien-
ists recommended a TLV of 1 mg/M^ and 0.5 mg/M^ for biphenyls of 42 and
54% chlorine content respectively^! )t
After 40 years of production Monsanto, under pressure from scientists
and government, voluntarily initiated a PCBs policy of sale only to
customers of closed system use. This largely stopped sales for such
applications as paints, papers, plasticizers, sealants, and adhesives,
leaving only controlled use in electrical capacitors and transformers.
Spent capacitors and discarded transformer oils could be returned to
Monsanto for incineration.
On July 16, 1973 EPA and the FDA, in response to a report by CEQ and OST
recommending discontinuation of the use of PCBs in all but electrical
uses, announced policies concerning PCBs in industrial effluents and
tolerable limits in foods and its related packagingC158)< EPA placed
PCBs on the toxic substances list as defined in the Federal Water Pol-
lution Control Act of 1972; it also announced it would limit discharge
of industrial wastes such that PCBs would not exceed 0.01 ppb in rivers
and lakes. The FDA simultaneously announced a temporary tolerance
level of 2.5 ppm in milk and dairy products, 5 ppm in poultry, and 0.2
ppm in baby foods. Also it banned the use of PCBs in plants manufac-
turing or storing animal feeds, foods, and food packaging.
PCBs have drawn sufficient concern to activate some international measures,
The World Health Organization in December 1972 announced it was studying
the health effects of PCBs and intends to issue environmental health
criteria( ). In February 1973 the Organization for Economic Cooper-
ation and Development representing Japan, Australia and the industrial
nations of Western Europe and North America adopted a directive calling
upon its members to limit the use of PCBs to transformers, capacitors,
heat transfer fluids (in other than food, drug, and feed operations),
53
-------
and hydraulic fluids for mining equipment; they also called for control
of the manufacture, import, and export of product containing PCBs' '.
5.11 Review of G.C. Analytical Procedures for Halocarbons
The literature that exists for gas chromatographic halocarbon analysis
deals primarily with separations at relatively high concentrations
characteristic of quality control production requirements. Reed^ J
evaluated the resolving power of a number of stationary liquid phases
for fluorocarbons and reported the ethyl ester of Kel-F acid 8114 to be
the most effective, Dresdner et al.(l62) used the same packing to
analyze 6- and 12- carbon fluorocarbon derivatives of SFg. Ellis et al.
(163) us«d a packing of Kel-F oil on Fluon powder to separate reactive
inorganic compounds such as GIF, HF, Cl2, C1F3, UF6 and Br2- Mixtures
of nitrogen trifluoride with CF4 were investigated by Nachbaur and
Engelbecht (.164) using moist silica gel with hydrogen as carrier gas
at 0°C. Campbell and Gudzinowicz (1°5) used diisodecyl phthalate and
Kel-F No. 3 oil on Chromosorb W with 1/4 inch columns of varying
lengths to measure several fluorocarbons and sulfur fluoride compounds.
demons and Altshuller(l66) determined the relative responses of an
EC detector to a number of halocarbons and sulfur halides using a Baymal
or an SF 96 column for separation. Greene and Wachi(167) used CH2=CHC02
(CF2-CF2)3H coated on Chromosorb W for the separation of several low
molecular weight fluorocarbons and Lysyj and Newton (168) found several
packing of halogenated polymers plasticized with halogenated oils useful
for the separation of some halocarbons. Williams and Umstead(169) used
Porapak Q and S in an on-line column concentrating device with temperature
programming to separate several halocarbons. Low concentration mixtures
(10 to 20 ppb) were concentrated to ppm levels (1 to 2 ppm) and then
measured by GC using a microcoulometric detector. This detector was pre-
ferred to the electron capture (EC) detector because of the latter 's
narrow linear range, lack of sensitivity to some compounds, and extreme
sensitivity to others.
Priestley et al. developed a GC method for phosgene using an EC
detector. An aluminum column packed with didecyl phthalate on GC 22
Super Support provided good separation. Jeltes et al.(-^l) used a similar
packing (diisodecyl phthalate) to determine phosgene and dichloroacetylene
in air at their subthreshold limit values.
Analytical methods for SF6 and a limited number of halocarbons have been
developed for meteorological tracer studies(20,172,173,174) _ Lovelock
(175) measured background concentration of ambient CClsF, CC14 and CHgl
over the North and South Atlantic using a silicone stationary phase and
a novel coulometric system. Hester et al.(176) used a stainless steel
column containing Na2S04 on Porosil A and temperature programming to
measure CCl^F and CC12F2 i-n L°s Angeles. The presence of atmospheric
CC14, trichloroethylene and tetrachloroethylene was reported. However,
based on the chromatogram presented, better resolution of components
would be necessary for accurate analysis.
54
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SECTION IV
EXPERIMENTAL
4.1 Reactors
4.1.1 Teflon bags and Myler bags -
4 ft. x 4 ft., five mil FEP Teflon bags were manufactured in the laboratory.
Teflon film in roll form from DuPont was folded over and heat bonded on
three edges with a 1/4" wide thermal-impulse heat sealer from Vertrod,
Inc. of Brooklyn, N.Y. Some compounds (C114, 111T, F-ll, F-12, F113)
were found to permeate through Teflon. In these cases identical bags
of Myler were manufactured and used. These compounds did not permeate
through Myler. The volume of the bags when filled was 200 liters (Teflon)
or 144 liters (Myler). Both bags were equipped with two ground-glass
ports for filling and sampling, and both were transparent to the wave-
lengths of the incident light. Irradiations were performed by exposing
the aluminized Myler-backed Teflon bags to a bank of 24 48" rapid start
40-watt fluorescent lamps. Mounted side by side the lamps formed a wall
52" x 52". The bank was enclosed by 4 walls (extending 22" outward),
which along with the reflectors of the lamp fixtures were covered with
reflective aluminized Myler. Thus the side opposite the light bank was
open in order to accomodate the Teflon bag. Four small electric fans were
mounted in the walls in order to ventilate excessive heat generated by
the lamps.
The use of 3 different types of fluorescent lamps permitted the approximation
of the quality and quantity of middle and near ultraviolet radiation found
in the lower troposphere. Twelve G.E. F40 BLB, 6 G.E. F40 BL and 6
Westinghouse FS40 fluorescent sunlamps were evenly distributed across tLe
24 mounts. The F40 BLB and F40 BL lamps emit most of the UV energy around
360 nanometers. The middle ultraviolet was supplied by the FS40 sunlamps
which emit their peak energy around 310 nanometers. Kj for Teflon bag
irradiations was 0.39 min.~ , and for Myler bags was 0.28 min.'l, due to
the lower transmission of Myler. Figure 5 shows the comparative spectral
distributions for this array and tropospheric radiation.
4.1.2 Pyrex Reactor -
A 72-liter Pyrex reactor was used in some experiments when long irradiation
times were desired. It was manufactured by Ace Glass of Vineland, N.J. and
was equipped with a thermometer, as well as 3 ports for sampling, filling
and venting. An all-glass, magnetic stirrer was enclosed in the sphere to
promote mixing. The reactor was used in an air conditioned irradiation
chamber equipped with 24 Westinghouse black light bulbs provided by the
Environmental Protection Agency. KI was 0.3 rain. .
4.1.3 Quartz Reactor -
A one-liter Hanovia photochemical reactor was used for studies in the U.V.
55
-------
TUFE 5
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56
-------
The reactor had a 450 watt high pressure mercury vapor lamp (>2200A).
4.1.4 Light Intensity Measurements -
The photolysis of nitrogen dioxide in air was used to measure the light
intensity generated by the lamp array used in bag irradiations, and also
for the array used with the 72 liter reactor. This measurement was made
at 30°C at zero relative humidity and was performed before, after, and
in the middle of all the experimental runs.
A 50 pphm concentration of nitrogen dioxide was produced in the bags
utilizing the nitrogen dioxide permeation tubes (see Section 4.3). After
ten minutes of irradiation, the steady-state concentrations of nitric
oxide, nitrogen dioxide and ozone were measured. With Kj as the reaction
rate for the reaction
N02 + hv kl NO + 0
>
in excess oxygen,
K = (25.5 ppnT1 min"1) [NO][0$]
1 [N02]
NO, 0, and N02 concentrations were measured at steady state, and Kj
calculated.
4.2 Preparation of Samples for Irradiation
4.2.1 Materials -
All chemicals were analytical reagent grade, or otherwise the best available.
Ultra-zero air manufactured by Matheson Gas Products of East Rutherford,
N.J. was used in all experiments except those in nitrogen. The air was
further purified by successive passage through an activated charcoal trap,
Ascarite, and molecular sieve. The air so treated had no detectable
halocarbon contaminants with the exception of CCljF which was present in
concentrations of less than lO'^v/v. Analysis showed less than 0.1 ppm
total hydrocarbon (expressed as Cffy) and less than 10 ppb NOX. Pre-
purified nitrogen was manufactured by Matheson Gas Products and was further
purified by passing it through traps containing activated charcoal, anhydrous
calcium sulfate and Ascarite.
The hydrocarbon fuel used in the photochemical experiments was a mixture
of paraffin, olefin and aromatic hydrocarbons. It was designated "Fuel A"
and was obtained from Exxon Research, Inc.
Its composition was 59.74% paraffin, 13.26% olefin, and 20.98% aromatic
hydrocarbons. Detergents and additives, mostly long-chain amino com-
pounds made up the remainder. This fuel mixture was used as a source of
free radicals in the simulations. Sufficient hydrocarbon was introduced
to the chamber, using a 10 microliter syringe to achieve 1 ppm.
57
-------
NC>2 was obtained from permeation tubes manufactured by Analytical Indus-
trial Development Co., West Chester, Pa. Allowing one week to obtain
thermal equilibrium the tubes were weighed on a Sartorius Analytical
Micro-balance and the weight plotted against time. The average leak-
rate for nitrogen dioxide tubes was 30,000 nanograms/minute.
4.2.2 Humidity -
50% relative humidity was used in all experiments. It was achieved by
injecting the calculated volume of distilled water into the filling gas
stream.
4.2.3 Temperature -
A tele-thermometer manufactured by Yellow Springs Instrument Company was
used to measure the bag chamber temperature. A probe was inserted into
the bag through one of the glass ports. A maximum temperature of 35°C
was reached after one hour of irradiation and the temperature remained
relatively constant thereafter.
4.2.4 Chamber Filling Procedure -
The Teflon or Myler bags were connected via a ground glass port to the
tank of ultra-zero air or N2 with its train of purifying traps already
described. A rotameter (Ace Glass #4-15-2) served to monitor the flow-
rate which was generally about 5 liters rnin."^. The test substrate, if
a liquid, was added to the inner surface of the port, so that the diluent
gas passed over it to allow evaporation into the bag.
In other cases, double dilution was used to achieve the desired concen-
trations, or permeation tubes were used. Double dilution involved
injecting a measured quantity of the substrate into a clean four-liter
flask through a. rubber septum. The flask had two stainless steel balls
to facilitate mixing. After allowing five minutes for evaporation and
diffusion a 5 cc aliquot was extracted and injected directly into the
bag chamber. Permeation tubes are described in Section 4.3.
Nitrogen dioxide was then added by attaching a Teflon line from the per-
meation apparatus to the bag. The time required for the desired NC>2
concentration to be achieved in the bag (50 pphm) was determined by
calculation from the known lead-rate of the permeation tube.
The bag was then mixed by slapping the sides of the bag 30 times. The
experiments conducted in bags employed no stirring except this initial
mixing of the chamber contents before irradiation. The initial concen-
trations of ozone, nitrogen oxides, nitric oxide, temperature and the
concentration of the substrate were determined (see Section 4.4). The
chamber was then placed in front of the light source. Before the
irradiation was started, the lamp array was turned on for approximately
30 minutes to allow the lamps to attain constant temperature and thus
constant output. Aluminized Myler was attached at the back of the bag
58
-------
chamber for use as a reflector. Measurements of substrate and product
concentrations were then made at 15 minutes, 30 minutes, 60 minutes and
at approximately one hour intervals thereafter. In the case of reactive
substrates irradiations usually lasted until a decay of one order of
magnitude was observed. For very long irradiations the nitrogen dioxide
was replenished daily to 50 pphm.
The 72 liter reactor, being a rigid chamber, required special filling
techniques. Cleaning was achieved by flushing 10 times the volume of
air through the reactor. Double dilution of substrate was employed.
After injection, the stirrer was switched on for 5 minutes. Nitrogen
dioxide replenishment was achieved by a 200 cc injection of a concen-
trated nitrogen dioxide mixture derived from a low flow off the per-
meation tubes into a 1 liter flask. The concentration of 50 pphm was
obtained empirically as the small size and rigid nature of the chamber
precluded sampling during a run.
Sampling was accomplished by attaching a Teflon line from the appropriate
instrument to the bag utilizing ground-glass ball joints for connections
and a stopcock to prevent back flow. Gas chromatograph samples were
withdrawn through a rubber septum placed in one port by an all glass
syringe and immediately injected into the gas chromatograph.
The bags were flushed twice with ultra pure air between each run. Each
experiment was duplicated.
4.3 Permeation Tubes
Permeation tubes were calibrated and used to provide known concentrations
of several of the compounds studied (F-ll, CC14, PCE, TCE, CH3I, F113 and
vinyl chloride) and of N02 and phosgene. The permeation tubes were main-
tained at 30.0 ± 0.1°C in a water-jacketed mixing chamber except where
noted in Table 4. Constant leak rates were usually established within 1
to 2 weeks of conditioning, as indicated by constancy in the rate of
weight loss when diluent air at 50 ml. min.~l was passed over them.
Table 4 shows the weight loss of the permeation tubes for CH3I, CC14,
CC12F2, CC13F, CH2C1F, CH3C1, CH2C1CH2C1, CC12, FCC1F2, CC12CHC1,
CH2CHC1, COC12. Tubes were weighed on an average about 3 times per week
using a semi-micro balance.
4.4 Laboratory Analyses
4.4.1 Gas Chromatograph -
The concentrations of the halocarbons were measured by direct injection
into a gas chromatograph. Two GC units were used in this study which are
referred to as systems A and B. System A was a temperature programmable
Fisher-Victoreen 4400 series GC and was equipped with a 15 mCi °^Ni EC
detector. Full scale deflection (FSD) on a 1 mv recorder was 10~1^A. It
was operated in the pulse mode (500 ysecs). Prepurified nitrogen was
used as a carrier gas. A flame ionization detector was used for vinyl
59
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chloride with helium as the carrier gas. System B was an isothermal unit
equipped with two identical 15 mCi Ni EC detectors in series. The
electrometer of each detector was connected to one channel of a Soltec
dual pen recorder. At maximum sensitivity FSD on a 1 mv scale was
3 x 10~^^A. The instrument was operated in a pulse mode (400 ysecs) .
The operating characteristics of this special purpose instrument have
been described in
Since water vapor would mask important peaks, a 2 in. x 14 in. o.d.
Teflon tube filled with Ascarite was routinely inserted (except for
phosgene analysis) between the EC detector and all columns used. For
Porapak Q above 80°C, the Ascarite was not very effective; however,
resolution was adequate. The Ascarite was replaced when the water peak
began to appear. All -glass syringes were used since Teflon gas-tight
syringes exhibited a continuous bleed of unknown interfering compounds.
4.4.2 Nitrogen Oxides -
Nitrogen oxides (N + N02) , nitric oxide and ozone were measured with an
Aerochem Model AA - III Chemiluminescence Monitor. The instrument is
based on the chemiluminescent reaction of nitric oxide with ozone. The
analysis for ozone required a constant source of pure nitric oxide where-
as the constant source of ozone for nitric oxide analysis was only dried
room air passed through an enclosed ozonator. The analyzer was specific
for ozone and nitric oxide; nitrogen dioxide was reduced to nitric oxide
in the instrument by passing the sample stream over a heated catalyst and
read as total nitrogen oxides. The monitor sampled at 2.4 liters per
minute with a 95% response time in 30 seconds.
Calibration was achieved by feeding a known concentration of nitrogen
dioxide in air from the permeation tubes and adjusting the potentiomete-
until the read-out galvanometer was correct. Nitric oxide was automatically
calibrated as well due to the 99% efficiency of the catalyst. The Ozone
mode was calibrated by performing an internal gas phase titration described
in the instrument manual. Double checking was done periodically using
both span gases of nitric oxide in nitrogen manufactured by MG Scientific
of Kearny, N.J. and the standard Saltzman method^ ^ for nitrogen dioxide.
4.4.3 Ozone -
Ozone was measured with a Rem Chemiluminescent Ozone Monitor. This monitor
is based on the chemiluminescent reaction of ozone with ethylene gas. The
reaction produces an excited species whose light signal produced in the
photomultiplier is proportional to the ozone concentration. The sampling
flow-rate was one liter per minute. The monitor was calibrated by setting
the gain control to equal a concentration measured by the neutral buffered
potassium iodide standard method. A Rem Ozone Generator was used to
generate a steady-state concentration of ozone for calibration and for
ozone dark reaction studies.
61
-------
4.4.4 Light Scattering Aerosol -
This was determined using an MRI integrating nephelometer. Particles of
radii between about 0.1 and 10 microns are measured in terms of their
scattering coefficient 8 scat. This is empirically related to the mass
loading after subtracting the molecular scattering contribution.
4.4.5 Phenols -
Phenols were collected by absorbing them into sodium bicarbonate solution
and analyzing by the colorimetric procedure described by Braverman et al.
( ). The analysis was based on the reaction of diazotized p-aminodi-
methylaniline sulfate with phenolic compounds to form a dye which can be
determined colorimetrically. With the collection of a 10:l gas sample and
the use of 40 mm spectrophotometer absorption cells a sensitivity of about
0.1 ppb was realized.
4.4.6 Aldehydes -
Aldehydes were determined by the sulfoxylate method(179)f xhe method is
sensitive to aldehydes which will combine with sodium bisulfite to form a
non-volatile addition product. The bisulfite is later released and
determined to quantitate the aldehydes present.
4.4.7 Phosgene -
Phosgene was analyzed with the Fisher Victoreen gas chromatograph. Sepa-
ration was achieved using 3 ft. of 1/4" o.d. aluminum packed with 30%
didecylphthalate on 100/120 mesh acid washed chromosorb P. A flow of
40 cc/minute nitrogen was maintained through the column along with 90 cc/
min of argon/methane (90/10) purge gas through the detector. The oven
temperature was operated at 30°C and the detector at 200°C. The detector
was operated in the pulse mode at 500 microseconds. The detector was
calibrated according to the pulse-flow coulometry method developed by us
and described in detail in Section 5.3. With a 10 cc sample the method
had a sensitivity of about 0.1 ppb. A calibration curve is shown in
Figure 6; standard phosgene concentrations were obtained using a per-
meation tube.
4.4.8 Chloride and Hydrogen Ion -
A sampling train consisting of a 5 mm i.d. teflon tube with an 18/7
socket joint, midget impingers, trap, rotameter, and vacuum pump was used
to collect HC1 from the bag samples. The gas was drawn through the im-
pingers which were each charged with 5 ml of distilled water, at the rate
of 1 liter min." for 10 minutes.
Chloride analysis was performed potentiometrically. A Corning Model 12
Research pH meter equipped with a Coleman 3-802 Ion Selective Electrode
and a Corning Double Junction Ceramic Reference Electrode was calibrated
62
-------
FIGURE 6
GAS CKROMATOGRAPHIC ELECTRON CAPTURE CALIBRATION
CURVE FOR PHOSGENE.
100
to
10
o
P-
CQ
§
1.0
0.01
-J l_
0.1
PP'A PHOS^E-VE X SA'-'PLE SIZE (MILLILITEHS)
1.0
63
-------
against standard Cl solutions at 25 C and were used to measure the Cl"
concentration in impinger samples. The EMF measured was compared against
the calibration curve of log (Cl~) vs EMF and the concentration obtained
(see Figure 7). The concentration of Cl" in the sampled gas was cal-
culated by the equation:
Concentration Cl" CV/V, ppm) = L X VL . 24.4 liters . ,n6
V mole
where L = choride ion concentration in the liquid analyzed moles/liter
VL = 1Q~2 liter = volume of collection water
Vp. = 10 liters = volume of air sampled
o
After the chloride analysis was completed the pH of the same samples were
determined. The same pH meter was employed with a Corning Combination pH
Electrode. The electrode was standardized against a phosphate buffer. A
similar calculation was performed in order to obtain the H+ concentration
in the air sampled:
H+ cone in ppm = [H+] x Vj . 24.4 1 . 1Q6
Vg mole
4.5 Field Analysis - Mobile Laboratory
A Winnebago motor home equipped with the required aerometric monitoring
instruments was used for on-site field studies initiated after May 1974.
The gas chromatograph equipped with a 10 mCi °^Ni EC detector and a flame
ionization detector (system A) was used for halocarbon analysis. The
other instruments used in this study were the Aerochem NO-NOx-Oj analyzer,
the REM ozone analyzer and the MRI integrating nephelometer.
Prior to May 1974, grab samples from various locations were analyzed in
the New Brunswick, N.J. laboratory. One hundred-mi all-glass syringes
each fitted with a three-way Luer-Lok valve and a septum were found most
convenient for this program and exhibited less than 10% halocarbon loss
over a 48 hour period. All-glass syringes were used for GC injections.
An all-glass manifold was used for ambient sampling of nitric oxides and
ozone.
64
-------
FIGURE 7
TYPICAL CHLORIDE ION ELECTRODE CALIBRATION CURVE.
PC
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CKLO°IDS ION ELECT^OIE POTENTIAL
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(MIILIVOLTS)
65
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SECTION V
GAS CHROMATOGRAPHIC ANALYTICAL PROCEDURES FOR TRACE LEVELS OF
HALOCARBONS AND SF6
5.1 Column Evaluation
The nine columns used in this study are given in Table 5 and are referred
to by their designated numbers. Unless otherwise indicated, data presented
here were obtained with System A. The optimal isothermal temperatures
for the separation of sub-ppb mixtures with air of SF^, CBrF?, CHC12F,
CC12F2, CC13F, CH3I, CC12F-CC1F2, CHC13, CH3-CC13, CC14, C2HC13, C2C14,
COC12 and CH2=CHC1 using the indicated columns, are given in Table 5.
For mixtures, optimal temperatures were chosen for adequate resolution
combined with minimum total elution time. For single compounds, this
elution time was limited to less than 10 minutes. Ppm concentrations of
CH3I, CCl2,F-CClF2, CHC13, CH3-CC13, CC14, C2HC1_ and C2C14 were generally
required for detection using column 1. The useful chromatograms of
ambient injections and synthetic ppb injections of SF6 and CBrF3 are
given in Figure 8. Figure 8 (C) also demonstrates the effect of purge gas
on the EC response to a 19 ppb mixture of CC12F2. Figures 9 and 10 are
chromatograms of ambient air samples using Chromosil 310 (column 2) and
Carbowax 1500 (column 3), respectively, where the latter includes a
chromatogram of a ppb synthetic mixture of CH2=CHC1 using a flame
ionization detector (Figure 10B). The chromatograms of typical ambient
air samples using columns 5, 6 and 7 are shown in Figures 11, 12 and 13,
respectively, and the dual trace coulometric chromatogram obtained with
System B using a DC 200 packing (column 8) is given in Figure 14. For a
typically polluted day in Bayonne, N.J., ambient concentrations cal-
culated from peak height versus concentration calibration plots are
compared with corresponding coulometrically calculated (equation 3,
Section 5.2.1) concentrations in Table 6. The chromatogram obtained
with column 9 (System B) of a 0.4 ml injection of an irradiated simulated
tropospheric photochemical reaction mixture of 1 ppm C2Cl4 and 50 pphm
N02 (50% RH) demonstrating phosgene synthesis is shown in Figure 5.
At relatively high concentrations (10 to 10* times ambient) the compounds
of Table 5 exhibited good separation on Porapak Q (column 1) at the
indicated temperatures, suggesting its applicability for ambient halocarbon
analysis. Using the Fisher Victoreen GC (System A), chromatographs of
ambient air samples (New Brunswick, N.J.) at each of the determined
optimum temperatures, however, revealed that only CC12F2 and CC13F could
be measured with this column (Figures 8A and 8B). Figure 8 also demon-
strates, using CC12F2 as an example, a phenomenon common to all halocarbons
under study: the use of purge gas (10% methane/argon) reduced the sensi-
tivity of the EC detector to these compounds. This effect of purge gas is
attributed to a decrease in ionization efficiency with increasing flow rate
as predicted from equation 1, Section 5.2.1. The use of purge gas accordingly
would only seem to be indicated at relatively high solute concentrations.
66
-------
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Figure 8
Ul
CC*
COLUMN I, 130° C
CARRIER (N2)FLQWRATE = 60mfi/MIN
AMBIENT INJECTION SIZE = 10.0mfl
FSD 8XIO~11 A
B
COLUMN I , 80° C
CARRIERIN2) FLOWRATE = 40mJ2/MIN
AMBIENT INJECTION SIZE = 5.0m fl
FSD 8XIO~" A
CCA 2F2
0.28 PPB
I I
PURGE GAS FLOW RATE
O.Om£/MIN
90m£/MlN
J I
COLUMN I , 80° C
CARRIER (N2)FLOW RATE=40mC/MIN
PURGE GAS 10% CH4/Ar
SAMPLE SIZE 1.0 mS.
FSD 1.6 X IO~'OA
ppb CCS. 2 F2 MIXTURE
I I I I I
COLUMN I , 30°C
FSDL6XIO'IOA
0.5 PPB
CARRIER (N2) FLOW RATE = 40mfi/MIN
SAMPLE SIZE I.Om.2
FSD8XIO"nA
12.4 PPB
T
T
CBrF3
6 8
MINUTES
10
12
14
68
-------
Figure 9
6
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COLUMN 2 , 23° C
CARRIER (N 2 ) FLOW RATE=20mfi/MIN
AMBIENT INJECTION SIZE = 0.5mje
COLUMN 2 , 23° C
CARRIER (N2)FLOWRATE = 30mtf/MIN
AMBIENT INJECTION SIZE = 4.(
FSD 1.6 X 1C
CC23F,0.8ppb
,0.25ppb
L
I
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8 19 20
22
MINUTES
69
-------
Figure 10
8
CO
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H20
COLUMN 3,23°C
CARRIER GAS(N2) FLOW RATE = 20mfi/MIN
AMBIENT INJECTION SIZE= 5.0 m£
FSD 8XIO""A
6
MINUTES
8 10
B
COLUMN 3, 23° C
CARRIER (He ) FLOW RATE = 20mfi/MIN
SAMPLE SIZE 10.0 mi
FLAME IONIZATION DETECTOR
14
I 4
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x
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X
CH2=CHCC
FSD8XIO"I3A
22.1 PPB
K FS
, l\ 2
-^^
24
MINUTES
70
-------
Figure 11
8
CCS. 3F
0.14 PPB
COLUMN 5,-l9°C
(LIQUID N2 USED FOR CRYOGENIC OPER.)
CARRIER (N2) FLOW RATE = 25mf/MIN
AMBIENT INJECTION SIZE = 4.0 m8.
FSD 8XIO"MA
2 4
MINUTES
Se-
6
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COLUMN 5 , 23° C
CARRIER { N2) FLOW RATE = 40mJ?/ MIN
AMBIENT SAMPLE 6.0 mi?
FSD 8XIO~" A
CH3-CCJ23,O.I5ppb
/
^CC^Alppb Q2ppb
1 1 1 1 1 1
0 4 8 12
MINUTES
71
-------
CD
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72
-------
E
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10
-------
The identity of both CC"L^2 anc* CC1,F was confirmed early in our experi-
ments by retention data on Chromosil 310 and Carbowax 1500 (Figures 9 and
10). These columns also resolved 0014. Reference to a chromatogram
obtained on a similar air sample using DC 200 and System B (Figure 14)
clearly illustrates that CH3I, CC12F-CC1F2, CHC13, CH3CC13, CC14, C2HC13
and C2Cl4 suffered irreversible losses on the Porapak Q column. Since
synthetic mixtures of SF^ and CBrF3 could be easily detected above 0.001
ppb and 0.005 ppb respectively, we must conclude that background SF^ and
CBrF3 are present at levels below these values. This is supported, in
part, by Lovelock^ ) who reported background SF6 concentrations of 1.2 x
10~1* v/v. Based on production figures CHC12F is probably present below
10~12 V/Y also. This is substantially below the EC sensitivity for this
compound (15 ppb) ; however, because of its inertness CHC12F should not
suffer significant losses on Porapak Q. It should be noted, however,
that SF^ and CBrF3, by virtue of their low background concentrations,
inertness, excellent electron absorbing properties, and good resolution
on Porapak Q at ambient temperatures (Figure 8D) , can serve as good
tracers for dual source studies.
The recent evidence indicating industrial exposure of workers to vinyl
chloride as a possible cause of a rare form of liver cancer known as
angiosarcoma supported the inclusion of vinyl chloride measurements in
our ambient monitoring program. Since the flame ionization detector was
found to exhibit a greater sensitivity to vinyl chloride than did the EC
detector, GC conditions were optimized with the former. A sensitivity of
10 ppb was achieved. No detectable vinyl chloride was measured at Sea
Girt, N.J.; New Brunswick, N.J.; Sandy Hook, N.J. or New York City (45th
Street and Lexington Avenue) during short term monitoring efforts in
June 1974. In our current monitoring efforts vinyl chloride concen-
trations as high as 1500 ppb have been measured in Delaware City, Delaware
(intersection of State Road 448 and Route 72) .
Based on the excellent separation of air and CC1F2-CF3, CH3-CHF2,
CC12F-CC1F2, CHC1F2, CC13F and CHC12F at high concentrations on a di-2-
ethylhexyl sebacate packingC-^"^) this packing was evaluated for ambient
analysis at various temperatures for the compounds under study. Only
CC13F could be detected, indicating significant sorption of the other
compounds by this column. An SE 30 packing on the other hand, was quite
useful. With this packing, it was possible to separate, identify and
quantify CH3CC13, CC14 and C2Cl4 in the ambient air at room temperature
(Figure 11) . Additionally, CC13F could be properly separated from the
air peak by cryogenic operation at -19°C (Figure 11) .
DC 710 (column 6) facilitated the detection of an unknown ambient com-
pound which subsequently was confirmed as CH3I (Figure 12) . Additionally,
this column provided excellent resolution of CC13F.
In an ancillary study to determine ambient concentrations of peroxyacetyl
nitrate (PAN), Carbowax 400 was found to resolve ambient CC13F, CC14 and
C2Cl4 (Figure 13) . This finding was fortuitous because it allowed
additional confirmation of the compounds identified on the SE 30 column
74
-------
by virtue of the significantly different characteristics of the two
columns. The unknown peak on the chromatogram remains to be identified.
By far the most useful packing for ambient halocarbon analysis was DC 200
(column 8). With this column, excellent resolution was obtained for
CC13F, CH3I, CC12F-CC1F2, CHC13, CH3-CC13, CC14, C2HC13 and C2C14 (Figure
14). Additionally, the separations were achieved isothermally at room
temperature. Confirmation of compound identity was obtained by use of
the previously presented retention data and the novel application of
ionization efficiencies. In this, the calculated ionization efficiency
(equation 2, Section 5.2.1) of a compound, tentatively identified by
retention data, was compared with that of a standard mixture for equality.
The relatively close agreement between the calibrated and coulometrically
determined concentrations for CC13F, CH3I and CC14 (Table 6), and our
experience with coulometric errors as a function of ionization efficiencies,
indicate minimal sorption errors for the DC 200 column. For the other com-
pounds, either sorption or coulometric factors(H) must account for the
discrepancy. As is clear from Table 6, where better agreement is seen
with the short column as compared to the long column, a large percentage
of C2C14 error is attributable to column sorption.
Although the tropospheric synthesis of phosgene from halocarbons has been
posulated, we are aware of no previous detection of phosgene in simulated
smog reactions. As can be seen from Figure 15, phosgene is synthesized
during the irradiation of a mixture of N02 and C2C14 with simulated sun-
light (>2950A ki=0.4/min). Confirmation was obtained, using standard
phosgene mixtures, with retention data and ionization efficiency charac-
teristics. The discrepancy between the actual (0.28 ppm) and coulo-
metrically determined (0.1 ppm) phosgene concentration is due to sorption
or heterogeneous decomposition in the column. Promising research now
underway in this laboratory indicates that this sorption problem may be
empirically corrected for by inclusion of sorption kinetics in the
coulometric analysis.
DC 200 has been demonstrated to be an excellent packing for the separation
and EC determination of ambient CC13F, CH3I, CC12F-CC1F2, 003, CH3-CC13,
CC14, C2HC13 and C2C14 under isothermal (23°C) conditions. Porapak Q at
80°C and Chromosil 310 and Carbowax 1500 at 23°C were found acceptable for
ambient CC12F2 analysis. Porapak Q at room temperature is also a very
good column for the separation of SFg and CBrF3. Carbowax 1500 exhibited
excellent resolution of vinyl chloride at room temperature (23°C). Phos-
gene was found to suffer unacceptable irreversible losses in all columns
studied except didecyl phthalate. This column provided excellent reso-
lution at 23°C and sorption losses, even though substantial, were suf-
ficiently small to allow accurate analysis in the sub-ppb range with
frequent calibration.
On a typically polluted day all halocarbons under study could be measured
in the sub-ppb range. As will be reported later, CCl2F2, CC13F, CH3I,
CH3CC13, CC14 and C2C14 are ubiquitous at their measurable concentrations.
The other compounds may not always be present at their measurable
75
-------
CM
S3HDNI-1H9I3H
76
-------
1 Figure 15
FSD3XIO •
COCC2
0.28
PPM
COLUMN 9, 23°C
CARRIER (N2)
FLOW RATE=60m£/MIN
SAMPLE SIZE 0.4 mfi
3 X IO"'°A
12345
MINUTES
77
-------
Table 6A. Comparison of Coulometrically Calculated
Concentrations with calibration Standards
CC12F-
Compounds CCl-jF CH3I CCl2F CHCls
Cone, based on
Calibrations 0.56 0.05 0.80 0.06 0.83
(ppb's)
Coulometric
Cone, (ppb's) 0.42 0.04 0.38 0.02 0.16
lonization 0.75 0.67 0.30 0.20 0.30
Efficiency
Table 6B. Comparison of Coulometrically Calculated
Concentrations with Calibration Standards
Compounds CCl4 C2HC13 C2Cl4 C2C14* COCl2**
rf
Cone, based on
Calibrations 0.19 0.75 0.50 0.53 280.0
(ppb's)
Coulometric
Cone, (ppb's) 0.16 — 0.22 0.39 100.0
lonization 0.80 — 0.60 0.60 0.75
Efficiency
* This was determined on a 2 ft. long DC-200 column in
which C2C14 eluted in 13 minutes instead of 33.
** Phosgene is reported as measured in an irradiated mixture
at 1 ppm C2d4 in air with 50 pphm N02 (50% RH; KI =
0.4/min).
78
-------
concentrations depending upon localized emissions in the sampling area.
The use of purge gas at low concentrations of solute leads to reduced
sensitivity. This is attributable to a decrease in ionization efficiency
with increasing flow rate through the EC detector.
The utilization of the ionization efficiency as a confirmatory aid has been
demonstrated and further applied to confirm the identity of phosgene in a
simulated photochemical smog study and, accordingly, the probably tropo-
spheric synthesis of this compound. Because of its high ionization
efficiency (>75%) and instability, the latter making the preparation of
calibration mixtures extremely difficult, phosgene and similar reactive
strong electron absorbers are excellent candidates for coulometric
analysis. Accordingly, for such compounds, new packings and passivation
procedures are recommended along with a detailed analysis of sorption
kinetics.
The exquisite sensitivity of the EC detector to most of the compounds under
study and their wide spectrum of emission patterns and reactivities make
them excellent candidates for the study of complex atmospheric phenomena.
The methods and data presented here should find wide application in studies
involving point and multisource diffusion modeling, transport of large
scale air masses, photochemical smog and sea breeze modeling and strato-
spheric exchange.
5.2 Gas Phase Coulometry
Under optimum conditions of near 100% ionization in an electron capture
(EC) detector, strongly electron absorbing compounds produce a response
of such .gfeat sensitivity that femtogram (10~15g) analysis has been sug-
gested^^). In this study a coulometric gas chromatographic method of
analysis based on gas phase electron absorption has been evaluated and
used for measuring the ambient concentrations of several such compounds.
Developed by Lovelock" ^, this method is based on a 1:1 equivalency at
100% ionization between the number of solute molecules in a carrier stream
to the number of electrons absorbed by them in the EC detector. Accordingly,
the solute concentration can be calculated directly from the number of
electrons absorbed. At ionization efficiency of less than 100% the use of
two identical detectors in series enables one to determine the fractional
ionizations in the EC detector and thereby maintain the absolute nature
of analysis by correcting for the unionized molecules. A further modi-
fication of gas phase coulometry, for use with reactive compounds which
undergo decomposition on the gas chromatographic columns, has been
developed by us. We have termed it "Pulsed Flow Coulometry" and it will
be described in section 5.1.3.
When operated coulometrically at 100% ionization efficiency an EC detector
is at its maximum sensitivity since all solute molecules are "counted."
Since the method is absolute, sources of mixing and contamination errors
inherent in the preparation of extremely dilute calibration mixtures are
79
-------
precluded. An additional advantage of this mode of operation, particularly
in air chemistry studies, is the ability to determine the concentration of
an unknown compound and possibly deduce its identity from spatial and
temporal distributions of concentrations. Furthermore, rate of thermal
electron attachment is an excellent aid for confirmation of identity when
used in conjunction with retention data.
5.2.1 Theory -
The theory of gas phase coulometry is as follows: Consider two identical
EC detectors in series and let their signals be Xj and X2 in coulombs due
to an injection of a plug of Q molecules of compound AB. If p is the
fractional ionization of AB, at high ionization efficiencies and low solute
concentrations one can write:
Xi = pQ (1)
X2 = p CQ-pQ) (2)
p = 1 - X2/XX (3)
The gram moles of solute W in the EC detector are therefore from equations
1 and 3.
W = Xi _ (4)
96,500 (1 - *1)
xl
For an injection of V ml of sample at t °C, the volumetric mixing ration
CAB is:
CAB =
= 8'5 x 10"4 x
Vrl - X2A
1
With a recorder the signals Xj and X2 are simply the respective chroma-
togram areas in coulombs.
Ionization efficiencies can theoretically be predicted by considering
the EC detector as a stirred tank reactor in which solute molecules under
go pseudo-first-order reactions with electrons present in large excess.
For a carrier gas passing through the detector at constant flow rate F0
simple mass balance gives :
iCVd (6)
where K is the first order constant and Vd is the detector volume. Veri
fication of the model offers the possibility of determing optimum para-
meters for coulometric analysis.
80
-------
5.2.2 Results -
The ionization efficiencies of the compounds studied are listed in
Table 7. The functional relationship between peak height and concen-
tration for CC14 and CClsF are shown in Figure 16. The variations of
ionization efficiencies with flow rates are plotted for these compounds
over a range of flow rates in Figure 17 and extrapolated from one point
through the origin for the other appreciably ionizable compounds (>30%)
of Table 7, Concentrations determined coulometrically at intermediate
efficiencies of 45% for CC^F and 50% for CC14 are plotted against known
concentrations in Figure 18. Coulometric convergence to actual concen-
tration with increasing ionization efficiency is graphed for these two
compounds in Figure 19. A chromatogram for a typical ambient injection
in the New Brunswick area is given in Figure 20 and the corresponding
coulometrically determined concentrations are reported in Table 8. The
chromatograms for the substances exhibiting greater than coulometric
response as shown in Table 7 are given in Figure 21 and the effect of
varying concentrations on this response is shown in Figure 22 for trans -
CHC1:CHC1.
5.2.3 Discussion -
CC14 and CCljF were evaluated extensively to verify the linear model
assumed in equations 1 and 5. Figure 16 shows that the EC detector is
slightly nonlinear over the concentration range studied. With con-
venient sample size (3-10 ml), this range includes typical to three
times ambient concentrations. Clearly, the profiles approach linearity
at very low concentrations and would appear to limit accurate coulometric
analysis to this concentration regime. Reference to Figure 18 shows,
however, for at least CC14 and CClsF this is not the case. The concen-
trations calculated with equation 5 are in fact linear with the standard
concentrations. This unsuspected result was due to a decrease in ioni-
zation efficiency with increasing concentration. The practical con-
sideration here is that by operating coulometrically with two detectors,
errors due to non-linearity of a single EC detector were offset over the
entire concentration range studied. This behavior is attributed to a
complex functional relationship between response, concentration, and
ionization efficiency. Qualitatively, this can be explained by considering
a non-linear response, suggested by Figure 16, of the type
Xi = pQ - m p2Q2 (7)
where m is a constant. Factoring equation 7
Xi = pQ(l-mpQ) where [mpQ]«l (8)
Similarly, the signal in the second detector is
X2 - P (Q-PQ) (1 - mp (Q - pQ)) (9)
Defining a detector coulomb efficiency E as the fractional change in
81
-------
Table?. lonization Efficiencies3
Compound lonization Efficiency
CC14 0.90
SF6 0.85b
CC13F 0.84
CBrF2-CBrF2 0.70
CC12=CC12 0.69
I 0.63
CC12F2 0.33
CH3-CC13 0.20
CCl2F-CClF2 0.10
CHCl3 0.06
CHCl=CCl2 0.00
trans-CHCl=CHCl < 0.00
CH2=CC12 < 0.00
<0.00
a_ , . __ 2.
2
Carrier Flow Rate: 33 cm /min. Porapak Q Column
82
-------
Table 8. Ambient Concentrations of
Coulometrically Determined Compounds in the
New Brunswick, N. J., Area
Compound Concentration ppb
CC13F 0.37
CH3I 0.08
CH3-CC13 0.27a
CC14 0.17
CrIC -L^CC -L f)
0.12
aBased on an ionization efficiency of 20%. Not amenable
to analysis (Ionization efficiency = 0) .
Table 9. Effect of Flow Rate on Detector Responses
Responses and Efficiencies for CC^F and
Dector
Compound
CCl3F
CC13F
CH3I
CH0I
Plow r.ate,
ml/min
71
11
71
11
.0
.3
.0
.3
response
coulombs
x 10^-0
209
45
135
16
.1
.0
.0
.5
Ionization
Efficiency
0.73
1.00 (.approx.)
0.60
1.00(approx.)
83
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100
• Figure 22 - Effect of trans-CHC1=CHC1 concentration on the greater than
coulometric response.
10 20 30 40 50
CONCENTRATION I PPM X SAMPLE SIZE)
GO
90
-------
coulombs X^-X2/X^ as distinguished from the definition of p (fractional
change in molecules),
E = 1 - £ (10)
Xl
By substitution of equations 8 and 9 in 10, one obtains by linearization
for the condition [mpQ]«l,
E = p - mQp2 (1-p) (11)
Clearly, E must always be less than or equal to p. The latter approached
E at extremely low concentrations (Q-K)), under linear conditions (m-K)),
as p-KL or a combination thereof. Equations 8 and 11 show a reduction in
Xi and E respectively as Q increases. The linearity seen in Figure 18
is attributed therefore to these effects which help maintain Xj/E as a
relative constant and thereby permit an extension of the workable coulo-
metric concentration range. This linearity does not preclude, however,
large errors at low efficiencies. The coulometrically determined CClsF
and CC14 concentration range of Figure 18. As the ionization efficiencies
approach 100% these errors become negligible as shown in Figure 19. It
follows that accurate coulometric analysis requires conditions favoring
maximum ionization efficiencies. From our data it seems reasonable that
barring irreversible sorption phenomena accuracy greater than 95% should
be routinely obtained at ionization efficiencies greater than 90%. Errors
of 25% and less can be expected at ionization efficiencies exceeding 50%.
According to equation 6, a plot of the reciprocal efficiency versus flow
rate should be linear. This is confirmed for CClsF, CC14 and CF^I in
Figure 17. The linearity allows one to determine an efficiency at one
flow rate and extrapolate it to zero flow rate. This has been done foi
all appreciably ionizable compounds of Table 7. Practically, one can
determine with such plots the feasibility and possible errors of using
coulometry. A large slope indicates poor efficiency which can only be
compensated for by very low flow rates. A small slope allows flexibility
in the choice of flow rates. However, although one can theoretically
achieve 100% ionization efficiency for an ionizable compound by approaching
zero flow rate, such flow rates may be imcompatible with practical analysis.
Table 9 gives the detector response at two flow rates for CC13F and CH3l.
The responses are seen to fall off drastically at the lower flow rate,
making the coulometric response invalid. This phenomenon was attributed
primarily to increased sorption and spreading in the GC column, which
could more than offset the advantage obtained due to increased efficiency.
It follows that one can only reduce flow rates as long as GC response
remains constant or increases.
Figure 20 is a chromatogram obtained from an injection of 8 ml. of ambient
air in the New Brunswick, N.J. area. Identified peaks are CCl^F, CH$I,
CH3-CCl3, CC14, CHCl:CCl2, and CCl2:CCl2. The identity of these peaks
was not only established by retention data on different packings (SE30,
Porapak Q), but also was confirmed by their efficiency or ionization.
91
-------
This was found to be a highly reliable method of identification since it
is extremely unlikely that two compounds will have the same retention
times as well as equal ionization efficiencies. Table 8 shows the
ambient concentrations of some of the compounds identified in Figure 20.
Hitherto, it has been considered highly unlikely that a greater than
coulometric response (Xi>X2) could be encountered in practice(182)_ jn
this laboratory, CC12:CH2, trans-CHC1:CMC1 and CH2C12 were found to
yield greater than coulometric response as shown in Figure 21. Figure 22
shows that for trans-CHC1:CHC1 the response of the second detector is at
least 265% greater than that of the first detector. Since such compounds
are apparently rare, dual EC detectors in series can be used reliable
to confirm their identification. Further, an increase in sensitivity is
possible by using the second detector conventionally. Based on the
mechanism for electron attachment proposed by Wentworth^°^), the
greater than coulometric response is attributed to the products of
ionization having greater electron affinities than the reactants.
5.5 Pulse Flow Coulometry
In the course of this study it became evident that certain highly
electron absorbing compounds, which otherwise would be prime candidates
for analysis by gas phase coulometry, presented real analytical diffi-
culties because of their tendency to undergo decomposition on the gas
chromatographic column. An excellent example of this is phosgene, which
is one of the strongest known electron absorbers (comparable to carbon
tetrachloride).
Previous workers, as indicated below, have reported difficulties with
phosgene analysis due to losses on columns. For example, Priestly et
al.' J reported the electron capture (EC) gas chromatographic deter-
mination of phosgene in air. Separation was achieved using an aluminum
column packed with 30% didecyl phthalate coated on 100/120 mesh GC 22
Super Support. Dahlberg and Kihlman^^^) USed a similar GC procedure with
a stainless steel column packed with 20% DC 200 on Chromosorb W. The
column had to be treated with acetyl chloride to preclude unacceptable
phosgene losses. A sensitivity of 1 ppb was achieved by both groups.
Jeltes et al.™^' used an aluminum column packed with 30% diisodecyl
phthalate coated on 80/100 mesh Aero Pak for phosgene analysis and re-
ported a sensitivity of 0.2 ppm with the EC detector.
Because phosgene was expected to be an important product in some of our
photochemical studies, and because of its potential importance as an air
contaminant, efforts were made to improve on the analytical procedures
available.
The wet chemical procedures reported in the literature for the deter-
mination of phosgene vapor have been reviewed by Kolthoff et al.0-85)
and Jeltes et al.^lj. in general, they suffer the typical difficulties
associated with wet chemical procedures, namely, lack of specificity,
92
-------
interferences, losses in sampling lines and the requirement of large
samples precluding real time analysis. Furthermore, such methods are
often elaborate and require considerable experience for acceptable
accuracy.
Because of its reactivity and the problems associated with wet chemical
methods, phosgene analysis is best accomplished by on-site GC procedures.
The three columns used by Priestly et al.^70)^ Dahlberg and Kihlman(l84)
and Jeltes et al.^1'IJ^ based on our experience, require routine cali-
brations because of the extreme reactivity of phosgene which causes
variable column losses, depending on the column's history. Clearly, an
absolute method not requiring calibration or suffering from changing
column characteristics is desirable. Reported here is an extension of
absolute coulometric analysis which empirically corrects for column
sorption through the use of pulse flow kinetics. For brevity the method
is called pulse flow coulometry (PFC).
5.5.1 Theory -
The theory of PFC is as follows: In the absence of column losses, from
equation 1 above we can write
Ce (ppb) = Ci (ppb) - 8.5 x 1Q5 (275 + t) Xj
- ^ ) (12)
where C^ is the inlet sample concentration, Ce is the exit concentration
expressed as the analogous mixing ratio of eluting C to V ml of mixture,
and KI and ^2 are tne respective serial EC detector responses in
coulombs. Use of equation 12 leads to accurate analysis only at high
ionization efficiencies (1 - X2/Xj) provided no losses take place in
the GC column. Reactive compounds such as phosgene, however, exhibit
a significant loss in the GC column requiring the novel application of
PFC for absolute analysis. In this procedure a pulse of the reactive
constituent is studied by coulometric analysis of the exiting constituent
to obtain the decay kinetics in the column.
Consider a packed column of length L and cross-sectional area S through
which a carrier gas flowing at a velocity u sweeps a slug of a reactive
gas mixture, say phosgene in air. For conditions of no axial dispersion,
mass balance gives the instantaneous moles of phosgene traversing the
column in time and space as,
9C 9C 8C* , ^
at + u 9! * 3t~ = - W
where r is the rate of reaction of component C in the GC column, C* is the
absorbed moles of C in equilibrium with the vapor phase, t is the time,
and 1 is the length along the column axis. Assuming first order kinetics
and a linear adsorption isotherm one can write:
93
-------
(r) = kC (14)
C* = AC (15)
where k and A are the first order rate and adsorption constants respectively,
Inserting equations 14 and 15 into 13, one gets:
3C u 3 kC
_
3t 1+A 91 = - 1+A
For the present model the following boundary conditions (B.C.'s) are
valid:
at t = 0, C = 0 for 1 > 0 (i)
at , n+ r W (t) for t > 0
at 1 = 0 , C = { (ii)
0 for t < 0
The solution to equation 16 with B.C.'s (i) and (ii) is given by Singh
et al.(4 ):
C = W (t - +3 exp (- -) (17)
Weight (or moles) of C out of the column Ce(ppb) Jo •._•,
T MO.AI
JoW(t - L l1+AJ)dt = J°W (t) dt, (19)
Weight (or moles) of C in the column Ci(ppb) >
0 W(t)dt
(18)
Since
JoW(t - iL_
u
inserting equations 17 and 19 into 18, it follows that:
Ce kV
£7 = exp (-|Sk) = exp (--p^-) (20)
where Fo (uS) is the carrier gas flow rate and Vc (LS) is the column
volume. Equation 20 can be rewritten as:
kvc
In Ce = In C- - ^— ,01,
f Q V^iJ
5.5.2 Results -
Standard samples of phosgene were prepared at several concentrations (10
to 167 ppb), and each sample was injected into the GC at several known
94
-------
carrier gas flow rates. Equation 12 was used to determine the column
exit concentration at each flow rate. A plot of In Ce versus 1/FQ was
made for each standard sample, and the intercept (C-jJ was compared with
the standard concentration.
Individual dual EC chromatograms of CC13F, CC14 and CC12F-CC1F2 are shown
in Figure 23 for standard injections of air-halocarbon mixtures prepared
with their respective permeation tubes. A comparison of the standard
concentrations with the coulometrically determined concentrations for
these compounds at their indicated ionization efficiencies is given in
Table 10. Figure 24 is a plot of phosgene permeation tube weight versus
time. Presented in Figure 25 are chromatograms of replicate injections
of a standard phosgene mixture. Figure 26 shows semi-log plots of Ce
versus 1/FO for several standard concentrations of phosgene. A used and
a freshly conditioned column were employed in these experiments. A com-
parison of the standard concentrations with those determined by PFC is
presented in Table 11. Figures 27 and 28 show the variation of ionization
efficiency with flow rate and input sample concentration respectively.
Figure 29 is a plot of phosgene decay in a conditioned glass vessel in
the presence and absence of water vapor. Also shown in Figure 29 is the
dry phosgene decay profile in a ground-glass vessel having a large spe-
cific surface area.
5.3.3 Discussion-
Coulometric analysis can be used for the absolute determination of elec-
tron absorbing compounds which do not suffer column losses. This is
illustrated in Figure 23 and Table 10, the latter of which shows less than
a 15% error between standard ppb injections of CCI^F and CC1$ and their
coulometrically determined concentrations (equation 12). For compounds
exhibiting low ionization efficiencies, however, significant errors
can be expected as is demonstrated by the 82% error for CC12F-CC1F2
analysis at a 28% ionization efficiency.
The reactivity of phosgene suggested the use of a TFE Teflon Permeation
tube, as a primary standard, for the dynamic preparation of ppb phosgene-
air mixtures. After a few days of conditioning, the tube exhibited a
constant rate of weight loss (Figure 24) and provided an extremely
reliable phosgene output (Figure 25). Figure 25 further shows that
phosgene is an excellent candidate for coulometric analysis because of
its high ionization efficiency (>85%) at a convenient carrier flow rate.
The assumptions of a linear adsorption isotherm and first order kinetics
for phosgene - column reactivity are proven valid by Figure 26, where
semi-log plots of C versus 1/FO are shown to be linear (equation 21).
As developed in the theory, the intercepts of these lines are the inlet
phosgene concentrations (Ci). From Table 11, PFC concentrations are
seen to be in very good agreement with the standards (maximum error =
10.2%). . This is attributable to the use of ionization efficiencies
exceeding 75% and the inherent applicability of the assumed kinetic
model. Since the lines are parallel for a given column (Figure 26),
95
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Figure 29 - Phosgene Decay in a Conditioned Glass Vessel in Presence and
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once the slope of one of them is determined, analysis is possible by a
single injection at one flow rate. The inlet concentration C^ would
then be determined as the intercept of the parallel extrapolated line. It
is recommended, however, that the slope of an analysis line be routinely
checked. The rate constant k may vary gradually with column history as
is suggested, in the extreme case, by the different slopes of the lines
obtained for the used and the conditioned columns. In conventional GC
procedures a change in slope would correspond to a need for a new cali-
bration whereas in PFC, the analogous requirement would be the injection
of an unknown sample at a few flow rates. This is substantially simpler
than a rigorous calibration, particularly for reactive gases like phosgene.
Since there is evidence'-" >^^ ' that carrier flow rate and solute mass
are the major factors governing the ionization efficiency and hence the
accuracy of the analysis, the effects of these variables on phosgene
ionization efficiency were studied. A linear plotf ^ >182 ) of the
reciprocal ionization efficiency versus flow rate (Figure 27) shows that
an increase in ionization efficiency can readily be obtained by reducing
the carrier gas flow rate. At a given flow rate, however, Figure 28a
clearly demonstrates that the ionization efficiency is nearly independent
of the input solute mass up to about 0.1 ng. Beyond this the ionization
efficiency is seen to fall off, presumably because of the decreasing
ratio of electrons to solute molecules (Figure 28b) . Consequently, in
the regime where solute mass has an effect on the ionization efficiency
(>0.1 ng for phosgene), the latter can be increased by a reduction of
solute mass as well as of carrier gas flow rate.
Having established the validity of PFC for the absolute determination of
phosgene in air, the method was used for a preliminary study of phosgene
vapor stability in glass vessels. A 560 ml cylindrical glass vessel
(45 cm x 4 cm i.d.) was flushed with a 167 ppb mixture of phosgene in dry
air for 60 minutes and was then capped and allowed to stand at 23°C fox
subsequent analysis. Figure 29a shows the first order decay of phosgene
in this system. After three hours, water- saturated air was injected into
this vessel by syringe to bring the water vapor concentration to 785 ppm.
It is clear from Figure 29a that water had, at best, a minimal effect on
phosgene decay. Figure 29b shows a much faster decay of a phosgene-air
mixture in a vessel of much higher surface area (ground glass), indi-
cating that the role of water, if any, is manifested in heterogeneous
surface reactions. Further, in tropospheric photochemical smog simu-
lation experiments where phosgene was synthesized from C2C14 in 200 liter
Teflon bagsC ), it was found that phosgene concentrations remained
stable for periods exceeding 15 hours in the presence of 10,000 ppm of
water vapor. We attribute the absence of surface reaction here to the
passivation of the walls by photochemical reaction products (CClsCOCl,
COC12, HCOC1, CC14, Cl2, HC1, etc.). These results lend support to
the observations of Noweir et al.(186) that the importance of phosgene
decay by the gas phase phosgene - water reaction has been overemphasized
in the literature.
105
-------
SECTION VI
AMBIENT MEASUREMENTS
Reported here are aerometric halocarbon and SF£ data obtained during
several programmed field studies initiated since March 1973 at various
locations in the U.S. which should be representative of a wide gamut of
emission patterns and meteorological conditions. The compounds included
in these studies were CClsF, CC12F2, CH^CCl^, CC12CC12, CC14, CHC1CC12,
CH3I, SF6, CH2CHC1, CHC13 and CC12F-CC1F2. Of these, several were sub-
jected to simulated tropospheric stability studies. Compounds that were
found tropospherically stable were tested for possible stratospheric
reactivity. Correlations between ozone and halocarbon concentrations
monitored at a non-urban location were used for the first time to
demonstrate the feasibility and utility of tracing large-scale air
masses for elucidation of photochemical smog phenomena.
6.1 Results
Figure 30 shows the locations of the eight monitoring sites. Table 12
presents the complete data for levels of all the above compounds at
Seagirt, N.J.; New York, N.Y.; Sandy Hook, N.J.; Wilmington, Del.;
Baltimore, Md.; Wilmington, Ohio; and Whiteface Mountain, N.Y. for the
dates and times listed during summer and fall, 1974. Also given in
the table are data for Bayonne, N.J. taken from March to December, 1973
for all the compounds except SF6 and vinyl chloride for which analytical
procedures had not at that time been perfected. Dashes in this table
indicate that analysis of the compound was attempted, but the level was
less than the lower limit of detectability.
Table 13 lists the maximum, minimum and mean values observed for these
eleven compounds at all eight locations.
Table 14 shows representative levels in urban and rural areas, as well
as the halocarbon levels associated with an inversion.
Table 15 lists the daily mean and monitoring period mean ambient levels,
and their standard deviations expressed as a percent of the mean. Be-
cause of the relative infrequency of monitoring at Bayonne, daily mean
values are not listed for this location. Instead, monthly mean values
are given in the table. Also indicated in this table are the number
of observations at which detectable levels of the compounds were present.
It should be noted in using the mean values in Tables 13 and 15 that
only those measurements at which detectable levels of compound were
observed are included in the mean, in some cases, using this procedure
undoubtedly gives a mean value which is too high; however, including zero
readings in the mean would, especially in the case of the non-ubiquitous
compounds, result in some of the "mean" values being less than the
minimum detectable level.
Table 16 shows for each compound at each location the precentage of
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analyses where detectable levels were observed.
In Table 17 an attempt has been made to quantify the degree of variation
of the ambient levels of the eleven compounds measured. For each com-
pound, a "weighted average standard deviation" has been calculated, by
summing all the individual location standard deviations with weighting
factors proportional to the number of measurements taken at each
location. The resulting values are listed in order of increasing
"variability."
This data is presented graphically in Figures 31 to 77 to demonstrate
more clearly diurnal trends. Table 18 lists aerometric parameters for
all times and locations monitored. For purposes of later discussion,
some representative atmospheric measurements taken at four of the moni-
toring locations are graphed together in Figures 78 - 81. The key to
symbols and scale factors used in Figures 31 - 81 is given in Table 19.
6.2 Discussion
Because of the need for a halocarbon data base, initial studies were
concerned with characterizing the spatial and temporal halocarbon
distributions. At all locations and times indicated in Table 13, CClsF,
CC12F2, CH3-CC13 and CC14 were always easily measurable. Accordingly,
they have been classified as "ubiquitous" for discussion. The minimum
detectable concentrations for the analytical methods used to measure
the non-ubiquitous compounds studied are given in Tables 13 and 14.
6.2.1 "Ubiquitous" Halocarbons -
The minimum concentrations that have been observed for the ubiquitous
halocarbons, CC13F, CC12F2, CH3CCl3 and CCl^, since June 1973 at the
time and locations indicated in Table 13 are 0.046, 0.06, .03 and 0.05
respectively. It is interesting that the values for CClsF and CC14 are
comparable to the reported background concentrations for the two, of
0.05 and 0.06 ppb respectively(I0>13). The minimum CC12F2 concentration
of 0.06 is an order of magnitude less than the concentration used by Su
and Goldberg(lS) to calculate a 30 year residence time for CC12F2- Their
background CC12F2 concentration of 0.7 ppb is inconsistent with cumulative
world production data. The integrated worldwide CC12F2 production would
give an average background concentration of less than 0.1 ppb.
During extended periods of inversion the ambient concentrations of the
ubiquitous halocarbons may attain values as high as 100-500 times their
minimum concentrations. For example, the maximum concentrations ob-
served for CCl^lr, CC12F2, CH^CCl^ and CC14 during a three day inversion
in Bayonne, N.J. were 8.8, 47.0, 9.8 and 18.0 respectively (Table 13).
137
-------
Table 17. Halocarbons and SF6 in Order of Ambient
Variability - Expressed as Weighted
Average S.D. for Each Compound*.
CC14 50.6
F-12 50.7
F-ll 59.6
SF6 66.1
HIT 66.5
C2C14 68.7
F113 95.4
CH31 96.5
C2HC13 112
CHC13 129
Vinyl chloride: Insufficient data
*Calculated as .Z. S.D. x No. of determinations
No. of determinations
138
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The maximum, minimum and average data of Table 13 for all the halogenated
compounds are indicative of a complex variable emission pattern and a
strong dependency on meteorological factors and should be representative
of the lower and upper limits one can expect at typical urban and non-
urban locations. However, the Bayonne, N.J. data show that, at times,
meteorological conditions are probably more significant than local emission
patterns. At this location not only were many of the maximum but also
minimum halocarbon readings obtained. The frequency of minimum halocarbon
data in Bayonne, N.J. was understandably quite low when compared to rural
area data, for example the Whiteface Mountains. Representative levels of
halocarbons that one can expect in urban and rural areas are given in
Table 14. Also included in Table 14 are data indicating a dramatic
increase in halocarbon concentration as one enters an inversion layer
from aloft.
The apparent ubiquitous nature of CClsF, 00^2, CH3CC13 and CC14, the
probability of their near exclusive anthropogenic origin, their signi-
ficant emissions, their low solubility in water, and theoretical chemical
and biological inertness all suggested, early in our experiments, that
these compounds would continue to accumulate in the troposphere and upon
photolysis in the stratosphere play a significant role in stratospheric
chemistry. However, there were insufficient data on CClgF and CCl2?2
and no data on CC14 and CH3CC13 which rigorously precluded their reactions
with typical reactive tropospheric species such as OH- , 0-, RO', etc. The
tropospheric stability of these four compounds is unequivocally demonstrated
in Figures 82 and 83. In these experiments kj for NC>2 photolysis was
0.39/min. Accordingly, in terms of total energy input, our laboratory
irradiation times are equivalent to much longer tropospheric irradiation
times. Similar experiments were conducted in which characteristic reactive
hydrocarbons were also present. Again, with the exception of the small
initial decay attributable to surface adsorption, these halocarbons were
stable. Since NC>2 was replenished periodically (up to 50 pphm NC>2 ever/
24 hours) during the course of the irradiation, reactive tropospheric
intermediates were continuously available for reaction with the halo-
carbons .
While atmospheric CC13F, CC12F2 and CH3CC13 are clearly anthropogenic,
the occurrence of CC14 in the atmosphere can not be accounted for from
direct production emission data*- ' J. No valid explanation for the CCl^
budget is available to date. Promising research in this laboratory
supports the possibility of considerable atmospheric CClq formation by
the photodecomposition of chloroalkenes in the troposphere ( J.
6.2.2 "Non-Ubiquitous" Halocarbons -
€2*^14 was measured at concentrations exceeding 0.06 ppb at least 50% of
the time at all locations. In a short-term monitoring effort in New York
City, levels as high as 9.8 ppb were observed. Similarly, CHC1CC12 was
observed over 90% of the time at all urban locations (except Baltimore)
at concentrations exceeding 0.05 ppb. A large clean oceanic air mass
had moved in during the Baltimore monitoring, which probably accounts for
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the absence of detectable CHC1CC12 as well as the generally low levels of
all other halocarbons. At the non-urban locations, however, this compound
was undetectable (<0,05 ppb) about 70% of the time.
Since the sources of C2C14 and CHC1CC12 are located primarily in urban
areas, urban transport plays an important role in their distribution.
However, superimposed on this is the tropospheric reactivity of these
compounds. The tropospheric photochemical reactivity of CHC1CC12 has
already been reportedW. Figure 84 is a concentration-time profile for
the irradiation of a mixture of 0.8 ppm C2C14 and 50 pphm N02 in ultra-
aero air (R.H.=50%). 02014 is clearly quite reactive and would be
expected to undergo significant photochemical degradation during transport
to non-urban areas. Depending on the 02014 source strength, meteorological
dilution, sunlight intensity and the presence of other trace constituents,
02014 or its predominant product COC^ accordingly may or may not be
observed in the non-urban areas. As is clear from Figure 85 the percentage
conversions on a chlorine basis of 02014 to 00012 is about 60%. Recognizing
the need for a sufficient reaction time, we calculate that an ambient con-
centration of 10 ppb 02014 observed in New York City should lead to the
formation of 12 ppb COC12 (TLV=100 ppb). Although quantitative data on
the 00012 yield from CHC1CC12 is not yet available, COC12 has been
unequivocally confirmed to be a major product of the tropospheric photo-
oxidation of CHC1CC12. The gas-phase phosgene-water reaction is probably
insignificant as a removal mechanism for COC^^ ), although particulate
surface reactions and reactions of 00012 with other trace constituents,
say NH3, may be important. This clearly suggests the need for a program
of ambient measurements of 00012- The rapid ozone decay in Figure 85 is
probably due to a chlorine atom ozone chain reaction similar to the one
proposed for the stratospheric ozone destruction by fluorocarbons f •*• ).
The lower frequency of detection of CHC1CC12 is probably due to a smaller
source strength(1^8) and a higher tropospheric reactivity^. Additionally,
however, urban transport may not always manifest itself in increased ground
level halocarbon concentrations in non-urban areas due to the trapping of
urban halocarbons in inversions aloft. The inversion data of Table 14
illustrates this point clearly. All halocarbon concentrations measured
at this rural location at ground level were typically low. Halocarbon
concentrations above 5000 ft. were similarly characteristically low.
Within the inversion layer however (1500 ft), halocarbon concentrations
were many factors higher, say 35 for CCl^F, than their corresponding ground
level or above inversion values.
CH3I, SF6, CH2CHC1, CHOI,, and CC12F-CC1F2, the other halogenated compounds
of interest, were generally measurable less frequently than 02014 and
CHC1CC12 and, for the most part, traceable to nearby sources. The note-
worthy characteristic feature of tropospheric CH^I is its general non-
detectability at locations remote from the ocean (Table 14). Indeed,
Lovelock et al.^ •'found a 100 fold higher mean surface ocean concen-
tration than aerial concentration for this compound and suggested the
ocean as a source. The lack of detectability at sites remote from the
240
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ocean is likely due to tropospheric reactivity (Figure 85). It should
be noted that a ghost peak frequently appears at the same retention time
as CHjI when one uses a glass syringe repeatedly. This is tentatively
attributed to heterogeneous CH$I synthesis from unknown ambient precur-
sors which accumulate on the glass walls.
The ambient SF^ concentrations were typically below 0.001 ppb although
levels as high as 0.005 ppb were observed in New York City. The latter
is attributed to leaks from SF6-dielectric high-voltage transformers,
or the only other probable source of this compound, a nearby meteorolo-
gist conducting a tracer study.
The reported link between Angiosarcoma and vinyl chloride monomer (pending
TLV=1 ppm) prompted the inclusions of CH2CHC1 in the halocarbon ambient
monitoring program. CH2CHC1 could not be measured at all locations
(sensitivity >10 ppb) except at Delaware City, Delaware, where levels as
high as 1.5 ppm were observed. These readings were obtained downwind
from a complex of chemical plants. Experiments conducted in this lab-
oratory which will be reported in detail at a later date demonstrate
that CH2CHC1 is tropospherically reactive and hence will not accumulate
in the troposphere.
Both CHC13 and CC12F-CC1F2 were detectable at most urban locations at a
relatively low frequency (sensitivity >0.01 ppb). CC12F-CC1F2 exhibits
significant tropospheric stability (Figure 86) and should accumulate in
the troposphere. Since CHCls absorbs radiation only well below the
290 nm tropospheric cut-off (CHC13 Amax = 175 nm) it can only undergo
thermal tropospheric reactions. Its structure however, suggests minimal
tropospheric decay by such reactions.
6.2.5 Ambient Variability of Halocarbons -
Table 17 lists the eleven compounds measured in order of the "variability"
of the ambient values, as defined in section 6.1. This quantity, though
extremely crude and statistically indefensible, nevertheless gives a
rough relative indication of the degree to which these compounds fluc-
tuate in their ambient levels. Interestingly, the compounds fall rather
neatly into two groups - a "low variability" group with values of 50-70%
and a "high variability" group, with values of 95-130%.
With the exception of C2C14, all the members of the first group are tro-
pospherically stable, which no doubt explains their relatively constant
values. CC14 is clearly the least variable compound, which is consistent
with the ideas already expressed regarding its possible sources.
The high variability of CH3l and C2HC13 is understandable in view of their
tropospheric reactivity, as well as the probable oceanic source of CH3l.
There were apparently high local sources of F113 at Bayonne, and of CHC13
at both Bayonne and Wilmington, Ohio, which contributed considerably to
the overall variability.
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6.2,4 Ambient Halocarbons as Tracer for Large Air Masses -
Because of their characteristic spatial and temporal distributions, wide
gradation of atmospheric reactivity, and amenability to ultra-trace
analysis, halocarbons are potentially useful for elucidating complex
atmospheric transport phenomena. Figure 87 is a diurnal concentration-
time profile for 03, light scattering aerosol (LSA) and several halo-
carbons obtained at Whiteface Mountains, New York, a relatively clean
air site (LSA 20 to 60 yg/M^). Ozone concentrations are seen to increase
gradually from 0800 to 2000 hours. The net increase of 20 ppb from 1100
hours to about 1600 hours can be attributed to either tropospheric syn-
thesis and/or transport. However, the ozone concaitrations increasing
until 2000 hours to a maximum of about 65 ppb can not be accounted for
by synthesis. Since the halocarbons generally originate in urban areas
their concomitant increase with ozone and particulates is an unequivocal
proof of urban transport. Stratospheric injection of ozone can be dis-
counted since this would be associated with deep vertical mixing and
hence, in the absence of significant simultaneous transport, a decrease
in halocarbon levels. Stratospheric injection occurring simultaneously
with transport can be discounted since CC14 would have to decrease, be-
cause of its known decreasing vertical concentrations profile ^ ^ com-
bined with its uniform distribution in urban and non-urban locations.
This is clearly demonstrated in Figure 88 where the levels of CC14 in
New York City and the Whiteface Mountains are remarkably similar. The
corresponding data for the other halocarbons shows their expected much
higher urban levels.
6.2.5 Seagirt monitoring - General comments -
At this location breezes were typically from the land in the morning and
from the sea in the afternoon. The pattern of NOX and 03 levels on
6/19/74 is consistent with this (Figure 78). In the absence of des-
tructive agents the ozone levels will remain through the night, an
observation made by others in some non-urban areas. With the onset of
the afternoon sea-breeze, all the fluorocarbon levels dropped. This is
especially noticeable with the ubiquitous compounds F-ll and HIT. The
drop in CC14 is less pronounced, consistent with its relative non-depen-
dence on location.
The data here demonstrate that caution needs to be exercised regarding
samples taken by the ocean, with on-shore breezes, as being representative
of background. At Seagirt, the halocarbon levels were clearly substantially
above background, concurrent with the Seabreeze. In other words, polluted
air is being recycled.
6.2.6 New York City monitoring - General comments -
The monitoring site was in the heart of mid-town Manhattan. Not unexpec-
tedly, the highest values for almost all compounds were registered here.
This is especially true of F-ll and €3014, which were each about an order
of magnitude higher here than at any other site. Local wind speed and
245
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direction are largely meaningless because of the proximity of tall
buildings - however the general wind direction over the city during the
monitoring period was northeast.
Extreme fluctuations in all the measured parameters are apparent, again
no doubt due to the peculiarities of the monitoring site. It is possible,
however, to discern the increases in nitrogen oxides during the morning
and evening rush hours. On 6/27/73 there is even a peaking to coincide
with the evening theatre traffic in this area (Figure 79).
6.6.7 Sandy Hook monitoring - General comments -
On all four days, exceptionally high maximum ozone levels were observed
(180 to 200 ppb). On the first day (7/2/74) this high 03 can be
attributed to high hydrocarbon and NOX emissions from inland being
transported over the monitoring site (Figure 80). The halocarbon levels,
which were characteristically higher than background, support this view.
At about 1300 hours the wind direction changed suddenly from north to
west. The air mass from the west was probably more contaminated with
ozone precursors, since there was a simultaneous dramatic increase in
C2C14 and, to a lesser extent, some of the other halocarbons.
NO was characteristically very low, which further supports the concept
of 03 formation upwind of the monitoring site, and is consistent with
the high 03 levels.
The monitoring period was just prior to a holiday weekend, with high
traffic density on the New Jersey Turnpike and other major roads, west
of the monitoring site. CO levels were seen to increase concurrently
with oxides of nitrogen, as would be expected.
Because of the possibility of analytical artefacts in the converter of
the NOX detector, care should be taken in considering the N02 data. The
N02 values could well include PAN, as is well known.
6.2.8 Wilmington, Delaware monitoring - General comments -
As far as halocarbons are concerned, this was a fairly "clean" site,
with F-ll and F-12 values as low as anywhere. Except for somewhat
higher CH3I values (due no doubt to proximity of the ocean), measurable
vinyl chloride, which has already been commented upon and lower CC14,
this location is not unlike the Whiteface Mountain site. 03 levels were,
however, exceedingly high, exceeding the 0.8 ppm standard. The air was
clearly contaminated, but in a "non-urban" fashion.
6.2.9 Baltimore, Maryland monitoring - General comments -
As stated earlier, a large clean air mass arrived from the ocean during
the monitoring period. The air was thus much cleaner than should normally
be expected in an urban location such as this. The very high NOX levels
248
-------
(Figure 81) are almost certainly due to the proximity of local sources.
The site was only a few hundred yards from the tunnel entrance of Inter-
state 95, which was vented by large fans. Also close was a D.O.T, main-
tenance building having very heavy truck traffic throughout the day. The
lack of wind during the night would permit oxides of nitrogen from these
local sources to remain in the vicinity.
6.2.10 Wilmington, Ohio monitoring - General comments -
Based on halocarbon levels, the air at this site was considered to be
mildly polluted. Surprisingly, CH^I was detectable on a few occasions,
despite its high tropospheric reactivity and the remoteness of the site
from the ocean. A local source is possible, such as the one which
apparently existed for CHC13. (CHC13 had the highest average value here
of any site, and the highest variability.)
The typical diurnal trend of 03 levels can be seen at this location, but
the absence of a diurnal pattern of NO and the high 03 concentrations even
in the evening, is indicative of transport.
A considerable amount of additional data was accumulated during this
monitoring period, as part of a collaborative effort with EPA. For
example, halocarbon levels were measured in aircraft samples taken above
and below an inversion line. A wide variation between the levels was
noted - for example, F-ll was 0.3 ppb above the inversion layer (6000 ft.)
and 8 ppb below it (4500 ft.). All this data has been supplied separately
to the EPA and will no doubt be incorporated in due course into a full
report on this monitoring effort.
6.2.11 Whiteface Mountain monitoring - General comments -
The evidence for large-scale urban transport to this location has already
been discussed (Section 6.2.4). Otherwise the site is the cleanest of
those studied, and the minimum halocarbon levels are probably typical of
background. As stated earlier, agreement with background values reported
by others is quite good.
6.2.12 Bayonne, New Jersey monitoring - General comments -
This was the first site at which regular halocarbon analyses were carried
out. Up till May, 1973 only F-ll, F-12, CC14 and CHsI were measured, but
after August 1973 the analytical procedures had been developed sufficiently
to permit measurement of F113, CHCls, HIT, TCE and PCE. The data here
were taken over a larger period and less frequently than at the other sites,
so it was not considered worthwhile to compute daily averages, as has been
done in Table 15 for the other locations. Monthly averages are given
instead.
This site showed perhaps the greatest variability of halocarbon levels. The
highest values of all compounds measured except TCE (where NYC had a slightly
higher value) were recorded here, as well as the lowest values of F-ll, CC14,
-------
TCE, CHsI, CHC13 and F113. The site is distinctly urban, but appropriate
meteorological conditions can evidently provide typical background levels,
on occasion.
-------
SECTION VII
PHOTOCHEMICAL REACTIVITY STUDIES
The compounds discussed in the following subsections (7.1 to 7.10) were
each subject to simulated tropospheric sunlight under a variety of con-
ditions . Those which showed no evidence of significant decay under
tropospheric conditions were further examined for their behavior under
simulated stratospheric sunlight conditions. The experimental details
have been described in Section 4.
7.1 F-ll, F-12 and F115
7.1.1 Results -
In all systems simulating tropospheric irradiation neither Freon-11,
Freon-12 nor Freon 113 reacted. Preliminary investigations had indicated
that these compounds would not react; as a result the following experi-
ments were pursued: Freon-11 and Freon-12 were irradiated in ultra zero
air at 50% relative humidity in Mylar bags for a period of 13000 minutes.
The system investigated contained Freon-11 and Freon-12 and N02- The
nitrogen dioxide concentration was monitored daily during the course of
this experiment and was replenished in an effort to maintain a concen-
tration of 50 pphm N02. Freon-11 and Freon-12 concentrations were also
monitored daily. The same system was investigated in a glass reactor,
and identical experimental procedures were followed. In both systems,
there was no net decomposition of either compound (indicated by no
decrease in concentration) as shown in Figure 82. A similar result was
obtained with F-113 (Figure 86). The aforementioned experimental scheme
was repeated with approximately 1 ppm hydrocarbon "A" in the system, and
the results obtained substantiated the previous findings (Figure 89).
Freon-11 was next introducted into the 72 liter Pyrex glass reactor. The
investigated mixture was composed of Freon-11, ultra-zero air at 50%
relative humidity and nitrogen dioxide. Nitrogen dioxide was replenished
daily to maintain 50 pphm, and Freon-11 concentration was monitored as a
function of irradiation time (almost 400 hours). The result of this
experiment (Figure 90) indicated no net decrease in the concentration of
F-ll indicating no photoreactivity of these compounds in a polluted
environment. A similar result was obtained with F-12.
Figure 91 shows the effect of simulated stratospheric photolysis on F-ll
and F-113. In these experiments the halocarbons in ppb binary mixture
with zero air were irradiated in the 1-liter quartz reactor by the 450
watt high-pressure mercury vapor lamp.
7.1.2 Discussion -
The results presented above conclusively establish that these compounds do
not react chemically in a polluted environment under the influence of
tropospheric sunlight. Their chemical inertness, and the fact that these
251
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Figure 91
HALOCARBON DECAY IN A I LITER
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compounds are transparent to wave lengths longer than 2900A^ •* serve to
substantiate our findings that photolysis of these compounds does not
occur in the troposphere. Molina et al.*- ^ have proposed stratospheric
photolytic dissociation of F-ll and F-12 to CFC12 and Cl, and to -CF2C1 +
•Cl, each reaction creating two odd electron species, a Cl atom and a free
radical. The ultimate proposed sink involves a catalytic chain reaction
leading to a net destruction of ozone and atomic oxygen as follows:
Cl + 03 ----- > CIO- + 02
CIO- + 0 ---- > Cl + 02
Figure 91 clearly establishes the stratosphere as a sink for these compounds,
From the slopes and a consideration of the average light intensity of
0.34 watts/cm used in our experiments - a level some 400 times the solar
irradiance at the top of the stratosphere ("^, the respective half-lives
of CClsF and CC12F-CC1F2 of 10 and 31.5 minutes correspond to minimum
stratospheric half-lives in hours of 69 and 217 hours.
With the expectation of 10-30 fold increase in concentration of freons in
the atmosphere, based on the assumption that present fluorocarbon pro-
duction and emission would continue, Molina and Rowland' ' calculated
atmospheric lifetimes in the range 40-150 years. Stratospheric photo-
dissociation of freons, it was proposed, would produce significant amounts
of chlorine atoms, and lead to destruction of atmospheric ozone by the
chain process above. In fact it was suggested C 190) that these compounds
in the stratosphere have already caused a 1% to 2% reduction in strato-
spheric ozone concentration, enough to cause an estimated 10,000 new cases
of skin cancer each year in the U.S. Reduction in the level of ozone
which heats the stratospheric air as it absorbs radiation could also
cause changes in glocal air circulation patterns and subsequent climatic
changes at the earths surface. The overall effects on man are difficult
to evaluate at this time. There could possibly be damage to the eyes,
sunburn ind aging of the skin.
It has, however, been pointed out' 19D that ozone concentrations less than
10% at atmospheric pressure and room temperature, the photosensitized
decomposition of 03 by chlorine is not a chain reaction, but has a quan-
tum yield of only 2. Further, the radicals can react not only with 0?
but also with other atmospheric constituents such as ^0 and N02 to form,
for example, HCL03, HCLO^ and HN03. Such competitive processes would,
of course spare the ozone from the full effects of the chlorine atoms
produced by u.v. photolysis of freons.
7.2 Perchloroethylene
7.2.1 Results -
In all irradiated systems investigated perchloroethylene reacted to 95%
completion. However, the rate of reaction was substantially different
in each system investigated. Hydrocarbons exerted an inhibiting effect
255
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on the reaction rates.
Perchloroethylene in Ultra Zero Air at 50% R.H, - In the photolysis of
perchloroethylene in ultra zero air at 50% relative humidity the compound
was 100% reacted in 1.5 - 2 hours, Figure 92, whereas in the same system
plus 1 ppm hydrocarbon the time needed for total reaction of this com-
pound was considerably longer (Figure 93) ,
Perchloroethylene in Nitrogen - Approximately 1 ppm was introduced in
the reaction chamber containing nitrogen. Figure 94 clearly indicates
the accelerated decomposition of perchloroethylene with a peak concen-
tration of ozone of 1 pphm 03. Perchloroethylene reacted to completion
in this system in about 7.5 hours yielding phosgene as a decomposition
product.
Perchloroethylene in Air at 50% R.H. plus N02 and Hydrocarbon "A" - Inves-
tigation of this reaction mixture indicated that the time needed for com-
plete decomposition of perchloroethylene was between 23 and 31 hours (4
runs shown in Figures 95 - 98).
Perchloroethylene and NC>2 in Ultra Zero Air at 50% R.H. - The time
needed for total reaction of perchloroethylene in this system was about
8 hours (Figure 84).
Dark Reactions in System Investigated - In all systems investigated dark
reactions were conducted to verify that these compounds would not react
unless supplied with the energy of the provided light source. This was
accomplished by preparing identical experimental mixtures and allowing
them to sit in the reaction chamber for the maximum time needed (indicated
by previous experimentation) to accomplish total decomposition of the
perchloroethylene if subjected to the light source. In all systems inves-
tigated concentrations of NO, NC>2, 0? and perchloroethylene were monitored
as a function of time. The data indicated no dark reaction.
Ozone Dark Reaction - In all experiments conducted it was observed that
at sometime in the course of the irradiation of the compound the ozone con-
centration would maximize followed by a sharp decrease in concentration.
This decrease in ozone concentration was always accompanied by a signifi-
cant acceleration in the rate of the perchloroethylene decomposition. It
was therefore decided to investigate the possibility of an ozone attack
on the perchloroethylene. This was accomplished by introducing a pre-
determined concentration of ozone (maximum ozone concentration observed
in experimentation) in air in the reaction chamber, and by monitoring the
ozone concentration as a function of time to determine the time needed
for thermal decay of ozone in the system. Approximately the same amount
of ozone was once again introduced into the reaction chamber with 1 ppm
perchloroethylene. The ozone and perchloroethylene concentrations were
monitored as a function of time. A decrease in the perchloroethylene
concentration and a decrease in the decay rate of ozone (indicated by
the slope) would indicate an "Ozone Attack" on the perchloroethylene,
256
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and would possibly indicate one path of the mechanism of the perchloro-
ethylene decay. The results, however (Figures 99 and 100) did not indicate
this as a path in the mechanism involved in the atmospheric fate of per-
chloroethylene.
Phosgene as a Product - In each system studied phosgene was identified
as a decomposition product. However, the rate of appearance of the com-
pound varied depending on the particular system that was being investi-
gated. Table 20 summarizes the systems investigated, and the phosgene
concentrations measured in each system as a function of irradiation time.
Phosgene as a Catalyst - It was observed in the system of air plus per-
chloroethylene that once the phosgene was detectable by the gas chromato-
graph, perchloroethylene concentration decreased very rapidly. Because
of this observation it was decided to investigate the catalytic effect
(if any) induced on this system by phosgene. This was accomplished by
introducing the maximum concentration of phosgene observed as a result
of the reaction of perchloroethylene reaction in this system (0.8 ppm) ,
and irradiating it for a period of time equal to the longest time needed
for perchloroethylene total decomposition. Phosgene and perchloro-
ethylene concentrations were monitored as function of irradiation time.
The results of this experiment, however, indicated that no catalytic
effect was exerted by the presence of the phosgene in the system, as
indicated by a lack of change in perchloroethylene concentration fol-
lowing 2 hours irradiation.
Formation of CC14 - CC14 has been observed by us as a product during our
longer runs in the simulated tropospheric irradiation of synthetic mix-
tures of C2C14 in air. In these and other long-term studies with other
compounds, N02 was replenished every 24 hours. This was necessary because,
in our experimental arrangement, the half-life of N02 was only about 6
hours (See Figure 101). Both COC12 and CC14 were identified by retention
data on two separate columns (5% SE30 and 30% didecyl phthalate for COC12,
and 5% SE30 and 10% DC200 for CC14) and by their electron attachment
properties( y>5>187 J, The results of these experiments are presented
in Table 21.
During the irradiations, an unexpected peak was tentatively identified as
CHC12COC1 and traces of CHClg were also observed. While quantitative
measurements of CCl^COCl were not feasible in these preliminary experi-
ments, it was identified as a product by the GC determination of its
isopropyl ester'1°4J>
7.2.2 Discussion -
The atmospheric fates of perchloroethylene were investigated in a variety
of systems. In each system investigated, perchloroethylene under the
influence of the light source employed was chemically reactive. The time
needed to accomplish a given decomposition of the perchloroethylene varied
significantly in repeated runs on the same system (see e.g. Figures 95 to
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OZONE DARK REACTIONS
OZONE PLUS C2CL4
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Figure 100
OZONE DARK REACTION;
DETERMINATION OF OZONE DECAY
DUPLICATE RUNS
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268
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98). Due to the highly irreproducible nature of this phenomenon, no
accurate explanation can be given at this time. However, Dusoleil et
al.(-29^ observed the same phenomenon in a series of experiments in which
they investigated rate constants in gas phase photochlorination of per-
chloroethylene and pentachloroethylene. Such erratic results they
attributed to "the time necessary for conditioning of the walls of the
reaction chamber"; accomplished by the occupation of the active sites
on its walls by "perhaps" polymeric chlorinated derivatives. As a re-
sult of this phenomenon they discarded a number of their initial exper-
imental results which were erratic in terms of total time needed for
100% decomposition of the compounds of interest.
None of the experimental results obtained by us were discarded, since the
primary objective of the derived experimental scheme was to indicate the
fates of these compounds in a polluted atmosphere. It was felt that the
variable nature of the decomposition kinetics were not as important to
an understanding of air pollution phenomenon as the observations that
C2C14 did decompose under simulated tropospheric conditions, and the
general behavior of the other reactants, intermediates and products.
Carbon Tetrachloride and Trichloroacetylchloride Formation - A maximum
concentration of 72 ppb carbon tetrachloride was measured as a decom-
position product of the induced perchloroethylene chemical reactions. This
was observed and measured in a mixture of air plus perchloroethylene
which had been subjected to the irradiating light source for 63 hours.
The same mixture was monitored for trichloroacetylchloride, and this com-
pound was positively identified as a decomposition product by the method
previously explained. CC14 concentrations continued to increase well
after all the C2C14 had reacted while the COC12 and CHC^COCl concen-
trations plateaued, At the same time CC^COCl continued to react suggesting
its role as the CC14 precursor. The overall reaction is probably initir t^ed
by photolysis of C2C14 followed by a chlorine sensitized photo-oxidation'" '.
From Table 21 on the average, C2C14 photodecomposition leads in 7 days to
the formation by weight of about 8% CCl4 and 70 to 85% of phosgene. Since
the present world-wide emissions of C2Cl^ can be estimated to be about 450
millions kg/yr.t1^)^ this conversion would result in an annual atmospheric
CC14 loading of 36 millions Kg/yr.
This source of CC14 alone can account, depending on the residence time, for
a significant fraction of the atmospheric CC14 budget. In any consideration
of the CC14 budget, this source must be considered along with others such as
direct industrial emission and possibly methane/chlorine reactions^™.
Our laboratory simulation studies indicate that CC14 is tropospherically
stable (see Section 7.4.1 below). Reactions such as CC14 + OH --> HOC1 +
CC13 or CC14 + 0 --> COC12 + C12 ' , are negligible under all tropo-
spheric conditions. Given the relative inertness of CC14, another
possible precursor of ozone destroying stratospheric chlorine atomsd )
presents itself in the form of CC14 as a secondary anthropogenic pollutant.
269
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It is significant too that a relatively innocuous material such as C2C14
(A.C.G.I.H. T.L.V, 100 ppm) can decompose in the atmosphere to produce
highly toxic materials such as phosgene (Proposed A.C.G.I.H. T.L.V. re-
vision 0.05 ppm).
Ozone Concentration - In all systems studied, it was observed that the
ozone concentration would maximize after some given time of exposure of
the system to the light source. The ozone concentration would then de-
crease with a simultaneous decrease in the concentration of perchloro-
ethylene in the system and an increase in the rate of phosgene produc-
tion (seeFigure 98) . In the system investigated in which hydrocarbon
"A" was incorporated it was observed that:
(a) Ozone concentration increased at a slower rate than in any of the
systems investigated, and decreased beyond its maximum more slowly.
(b) The reaction rate of perchloroethylene was substantially reduced.
(c) The time needed for 96 to 100% decomposition of perchloroethylene
was substantially increased.
(d) The observed ozone concentration was substantially greater than that
observed in any other system investigated.
Possible Reaction Mechanism for C2C14 - There are no documentations of
simulated smog chamber studies performed on perchloroethylene. However,
gas phase photochlorination studied performed using intermittent light
capable of inducing chemical reactions have been performed on mixtures
of perchloroethylene with Cl2 in airf^o) ^n which the following scheme
was proposed as the initiator in the observed chemical reactions of
perchloroethylene .
Cl2 + hv ----- > 2C1
C2C14 + Cl ----- >
C2Cl5 + C12 ----- > C2C1& + Cl
c2ci6 + ci ----- > c2cig + ci2
2C1 + M ----- > C12 + M
C2Clg + CL ----- > C2C14 + C12
2C2Cl5 ----- > C2C16 + C2C14
They also established the fact that reaction with impurities on the walls
or other entities (Cls, polymeric chlorinated derivatives needed to occupy
sites on the walls of the reaction chamber) were, within a highly experi-
mental precision, completely negligible. Although the experimental scheme
studied by this investigator did not involve introduction of chlorine in
the reaction chamber, the writer proposes a chlorine sensitized oxidation
of perchloroethylene, initiated somewhat differently from that proposed
by Dusoleil et al . , but including some of his proposed paths. The pro-
posed mechanism is as follows:
270
-------
C2C14 -_v_> C2CI^ + Cl
Cl + C2C14 ----- > C2Cl5
C2Cl5 + 02 ----- > C2C1502
C2ci5o2 ----- > c2cij; + o2
C2Clg + C2C1502 ----- > C2C1502C2C15
C2C15°2 + C2Cl502---> C2C1502C2C15 +
C2C1502 + C2Cl502 — -> 2C2C150' + 02
C2C150- ----- > CC13COC1 + Cl-
c2ci5o- ----- > coci2 + cci^ - >coci2 + ci
+ Cl- ----- > CC14
As is clearly indicated by the proposed mechanism there is a time lag
necessary for the initiation and propagation of the radical -chlorine,
radical -oxygen and radical -radical species. Figure 92 indicates a
time lag followed by a rapid decomposition of perchloroethylene. In the
systems investigated in which hydrocarbon is introduced there is an
extensive initial time lag before perchloroethylene begins to react
chemically, and this is followed by a slow reaction rate. This can be
explained by the fact that the very reactive olefin radicals not only
compete with the C2C14 for the C2Cl5 specie but also probably react
rapidly with the Cl chain carrier itself. This proposed path would be
followed until the olefinic specie from the hydrocarbons has been totally
reacted. Then the perchloroethylene would be attacked by the Cl specie,
hence the observed extended time lag in the hydrocarbon system. An
additional explanation could involve collisional de-excitation of the
C2C14* with olefin, resulting in a photosensitized dissociation of the
latter.
Step 8 in the proposed reaction scheme was substantiated by identifying
trichloroacetyl chloride with extended irradiation time of 24 hours. It
was observed that the trichloroacetylchloride had been completely decom-
posed and the concentration of carbon tetrachloride had increased con-
siderably. The overall reaction scheme may be presented to proceed as
follows by consecutive or side reactions:
C2C14 ----> COC12 + CC14 + CC13COC1
Blank runs of all systems investigated produced very low NO, N02 and 0^
readings and no carbon tetrachloride or trichloroacetylchloride.
7.5 Methyl Iodide
7.5.1 Results -
Figure 102 shows the stability of methyl iodide in the presence of ozone
which decays thermally while the methyl iodide concentration remains the
same.
271
-------
272
-------
Concentracion (ppb)
(iLdd)
273
-------
Figure 103 represents the photolysis of methyl iodide in a nitrogen
atmosphere. The photolysis is apparently first order as the straight line
plot (Figure 104), of the concentration vs. time shows. The half-life of
the methyl iodide in this experiment was 6.5 hours.
The photolysis of methyl iodide in the presence of oxygen, as pointed out
in the literature^46,49,50,51,194 J should be faster than in its absence due
to the fact that methyl radicals react with oxygen rather than recombine
with iodine to form methyl iodide. However, the Teflon bags used for runs
were permeable to the oxygen in the room air and therefore the half-life
of 7.6 hours (Figures 105 and 106) for the photolysis of methyl iodide in
the presence of air and water is in disagreement with neither the previous
value nor with other investigations. Figure 85 shows the photolysis of
013! with the formation of NO and 63 in air in presence of N02 and 50%
relative humidity, and Figure 107 shows the first-order decay of CH^I in
this system. Figure 108 shows the same measurements when 1 ppm of
hydrocarbon ("Fuel A") was present in addition.
7.3.2 Discussion -
From Figure 85 it is clear that the NC>2 half-life in air is much smaller
in presence of CH^I than in its absence (Figure 101). With CH^I present,
the NC>2 half-life was less than 2 hours and the NO reached a peak of
0.175 ppm, whereas in the blank the maximum nitric oxide value was 0.120
ppm. the initial concentrations of nitrogen dioxide were respectively
0.450 and 0.535 ppm. Note that the nitric oxide remained at a steady-
state concentration in the blank but was consumed in the methyl iodide
irradiation. From the literature&*• >^*' 1"5) nitric oxide has been shown
to be an extremely efficient scavenger for methyl radicals. The half-
life of methyl iodide is a fast 3.6 hours (Figure 107) due to the
inhibition by nitric oxide of the back reaction reforming methyl iodide.
Hydrocarbons seem to negate the influence of nitric oxide on the reaction
(Figure 108). The half-life of methyl iodide in this simulated urban air
was 5.5 hours.
The source of atmospheric methyl iodide has been postulated by Lovelock et
al.*- •'to be the oceans. Biological methylation of iodine by marine plants
and animals is a likely cog in the wheel of the iodine cycle. Another
possible way the sea can be a source for methyl iodide is by methyl radi-
cals in the air reacting with sea salts. To test this possibility, a mix-
ture of 12 ppm of potassium iodide and 1.0 ppm of acetone were irradiated
in the 72 liter pyrex reactor. Almost immediately a four-fold increase
in methyl iodide over the unirradiated blank to a maximum of 1.4 ppm of
methyl iodide was observed. This result agrees with experiments of Pitts
and Blacet ( 195) .
7.4 Carbon Tetrachloride
7.4.1 Results -
Figure 83 shows the reaction of carbon tetrachloride with nitrogen oxides,
274
-------
(qdd)
275
-------
- •„ Nitroeen on semi-log paper
4 Figure 104 - CH3I Photolysxs in Nitrogen,
4 . 6
Time Irradiated (Hrs.)
276
-------
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10 \
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Figure
semi-log paper
CH,I
278
-------
Figure 107 - CH3I * N02 in air with 50% R.H. on semi-log paper
12
279
16 Time Irradiated (Hrs.)
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(u>dd) ucnriEaauaoucrj o tueajouj
280
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ozone, atomic oxygen and simulated tropospheric sunlight in a Mylar bag.
The compound remained stable for more than 200 hours. The chamber was
replenished daily with nitrogen dioxide to bring the chamber to a 0.5
ppm concentration providing a constant supply of reactive constituents
for attacking the halocarbon.
Figure 109 depicts the stability of carbon tetrachloride from the attacks
of free radical hydrocarbons as well as ozone, atomic oxygen and nitrogen
oxides. Here too, nitrogen dioxide was replenished daily.
Having established the tropospheric stability of CCl^, its possible
stratospheric reactivity was investigated. Figure 90 includes its
decay under the influence of U.V. irradiation in ppb mixture in air.
Its half-life under our conditions of 5.5 minutes corresponds to a
minimum half-life in the stratosphere of 37 hours.
7.4.2 Discussion -
From the data presented above it can be inferred that carbon tetrachloride
will not decompose in the troposphere. It is clear, however, that €014
will play a role similar to that ascribed to CCljF and CC12F2 as precur-
sors of chlorine atoms potentially destructive to stratospheric ozone.
7.5 1:1:1 Trichloroethylene
7.5.1 Results -
Figure 109 shows the reaction of 1:1:l-trichloroethane with nitrogen
oxides, ozone, atomic oxygen, free radicals and simulated tropospheric
sunlight in a Mylar bag. The compound remained stable for more than
200 hours. The chamber was replenished daily with nitrogen dioxide to
bring the chamber to a 0.5 ppm concentration, providing a constant
supply of reactive constituents for attacking the halocarbon.
7.5.2 Discussion -
From these systems, it can be inferred that methyl chloroform will not
decompose in the troposphere. We have not tested its stability to strato-
spheric conditions.
7.6 Trichloroethylene
7.6.1 Results -
Figure 110 illustrates the stability of T.C.E. to thermal attack by
ozone. Figure 111 shows the effect of irradiating this compound with si-
mulated tropospheric sunlight in nitrogen. The effect of the presence of
air and 50% relative humidity are shown in Figure 112. With the addition
of N02 to the reaction mixture the results shown in Figure 113 were
obtained. Figure 114 shows the results when both N02 and hydrocarbon
281
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Figure 110
TRfCHLOROETHYLENE +
dark reaction
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("Fuel A") were present in the irradiation mixture.
7.6.2 Discussion -
T.C.E, is clearly stable to direct 03 attack - the ozone half-life was
unchanged whether T.C.E. was present or absent, and no discernable
decrease in initial T.C.E. concentration was apparent after almost 60
hours of exposure to 63.
In the system where N02 was employed, it was noticed that the time
required for 95% decomposition was much longer than in systems where
trichloroethylene was reacted without the nitrogen dioxide present. This
same phenomenon was observed by Wilson(21) when she studied the photo-
reactivity of trichloroethylene. It was noted by her that if the N02
concentration is several times that of trichloroethylene, the reaction
was completely inhibited and that the rate of decomposition is quite
dependent on the T.C.E./N02 ratio. As the ratio decreased, she noted
that all the measures of photoreactivity also decreased.
In the system where both nitrogen oxide and Fuel A were introduced, it
was noted that the hydrocarbon further inhibited the reaction. The
occurrence of a competitive reaction seems to be a good possibility. The
synergistic effects of one hydrocarbon to another are extremely hard to
evaluate, and since this report is not primarily concerned with reaction
kinetics the possibilities were not thoroughly investigated.
In all the systems investigated, phosgene was the major decomposition
product observed. By doing a chlorine balance it can be seen that at
the point of 95% decomposition approximately 30% of the budget is converted
to phosgene. Depending on the system studied the rate of formation of
phosgene varied;the quicker the trichloroethylene dissociated, the more
rapid the formation of phosgene. A proposed mechanism for the formation
of phosgene and the other decomposition products which is consistent with
the present observations is presented later in this section.
Chloroform was noted to form in all systems where light was employed, the
maximum concentrations observed varying from 8-15 ppb depending on the
system investigated. Again, the rate of formation was dependent on the
decomposition rate of T.C.E. Other products identified were dichloro-
acetyl chloride and HC1. Both of these products were identified quali-
tatively but the concentrations could not be quantified due to limita-
tions of procedures employed. In order to account for these products a
chlorination reaction is suspected.
The formation of ozone occurred in all the systems investigated where a
light source was employed. After the ozone concentration reached its
peak a steady decay took place. It was also noted that following the
peak the decomposition of trichloroethylene accelerated. The time for
the peak concentration to occur in identical systems varied considerably
as well as the magnitude of the peak concentrations observed.
288
-------
The following reaction mechanism is proposed in order to account for the
observed formation of the products mentioned. The mechanism involves a
chlorine-sensitized photo-oxidation of trichloroethylene. Huybrechts et
al. *- ' did work on the chlorine photosensitized oxidation of trichloro-
ethylene but in their case CL2 was added as a primary reactant, The
mechanism they proposed was discussed earlier in this report. This
mechanism could not account for the results observed in this study but
some of the reactions are included. The proposed mechanism for the
present study is as follows:
i) c2HCi3 -----> c2Hci2 + cr
2) Cl- + C2HC13 > C2HC14-
3) C2HC14 + 02 > C2HC1402'
4) C2HC14' + C2HCl402--> C2HC1402C2HC14
5) C2HC1402- + C2HCl402---> 2C2HC140' + 02
6) C2HC140- > CHC1COC12 + Cl-
7) C2HC140- > COC12 + CHC12'
8) CC13-CHC10- > CC13' + HC1 + CO
9) CC13' + %}2 > COC12 + Cl-
10) CHC12' + Cl- > CHC13
This mechanism accounts for the products and the time lags observed in
the various systems investigated. Time is required for the initial pro-
pagation of chlorine radicals, oxygen radicals and radical-radical species.
The time lag varied significantly from one system to another due to the
presence of N02 and fuel competing for the radicals being formed. The
time lag in the systems with nitrogen dioxide is substantially longer than
in the systems not incorporating N02. This is suspected to be due to
nitrogen dioxide competing for Cl• radical thereby increasing the time
required for trichloroethylene decomposition and product formation. Evi-
dence to this occurring is provided when the chlorine budgets in runs
incorporating N02 and those not incorporating N02 are compared. The products
observed in the runs not incorporating N02 account for approximately 41%
of the chlorine initially present, where in the systems with N02 added the
products observed only account for about 24% of the chlorine. A possible
competitive reaction follows:
Cl + N02 + M > C1N02 + M
C1N02 + 0 > C1NO + 02
If this reaction occurs it can account for the differences in the chlorine
budgets. Further evidence of this competition occurring is provided by
Figures 113 and 114. These figures demonstrate that as the N02 concentration
drops off the inhibition becomes less severe and the trichloroethylene
decomposes at a much more rapid rate. Similarly, the mechanism can account
for the increased time lag when Fuel A is added to the system. In this
case again a competition for free radicals is expected. The reactive
289
-------
olefin species provided by the hydrocarbon competes with the trichloro-
ethylene for the 02^1^' species as well as the Cl' chain carrier. Once
the olefinic species is reacted, then the photo-dissociation of trichloro-
ethylene proceeds at a much more rapid rate.
7.7 Ethylene Dichloride
7.7.1 Results -
Figure 115 shows the stability of dichloroethane to simulated tropospheric
conditions. No decomposition was observed over a period of almost 400
hours.
7.7.2 Discussion -
Ethylene dichoride is a saturated alkane, held together only by sigma
bonds. Extremely high energy of short wavelengths are necessary to pro-
mote the electrons.
In order for any compound to react photochemically, it must be capable of
absorbing light in the range of wavelengths employed. In this study the
wavelengths of light employed were predominantly in the visible region
of the spectra. In order to determine if the compound absorbs in the
region the spectrum covering the UV and visible portions was examined,
and it was found that -it was transparent to wavelengths greater than
2900A. Therefore, light present in the chamber or in the troposphere
could not cause the compound to react.
It is also evident from Figure 115 that dichloroethylene is also very
unreactive towards reactive photochemical intermediates. Tropospheric
reactions are thus unlikely to be a major sink for the compound.
7.8 Vinyl chloride
All experiments were performed using 200 liter Teflon bags filled as
previously described (Section 4.2.4). Vinyl chloride initially at
F*l - 10 ppm was irradiated with simulated sunlight under the following
conditions:
a) in ultra zero air at 50% relative humiditiy
b) in dry nitrogen
c) in ultra zero air at 50% relative humidity, with N02 present,
initially at 0.2 to 6 ppm
d) as for (c), with hydrocarbon present.
The dark reaction between vinyl chloride and ozone (~5 ppm) in ultra zero
air at 50% relative humidity was also studied.
7.8.1 Results and Discussion -
The effects of irradiating vinyl chloride in air at 50% relative humidity
290
-------
291
-------
are shown in Figure 116. For comparative purposes the corresponding
data for a nitrogen atmopshere are also shown. It can be seen that in
both cases the removal of vinyl chloride is rather slow, with only a
slightly greater rate in presence of oxygen. The removal rates cor-
respond to half-lives of approximately 17 hours in air and 30 hours in
N2.
Several experimental runs were performed in which N02 was present in the
vinyl chloride/air mixtures. Results of one such run are shown in
Figure 117. The initial vinyl chloride concentrations were varied between
1 and 10 ppm, and the vinyl chloride:N02 ratios were in the range 1.4 to
5.3. No significant differences were observed in the time for the vinyl
chloride concentration to fall by a factor of two - the average time was
3±0.6 hours. Ozone usually peaked at about this time. The other reaction
products identified in the V.C./N02 runs were NO, formaldehyde and HC1. NO
was usually very low, peaking soon after the start of the irradiation, and
falling rapidly to undetectable levels within about 3 hours. Formaldehyde
slowly built up over this period to 1 or 2 ppm. HC1 was not quantitatively
determined.
The effect of adding hydrocarbon to the V.C./N02/air system is shown in
Figure 118. Though the N02/V.C. ratio was very low, the vinyl chloride
still disappeared at about the usual rate. NO disappeared very rapidly,
accompanied by a rapid rise in Oj, and a steady build-up of formaldehyde.
Vinyl chloride in air reacts rapidly with ozone in the absence of light,
as shown in Figure 119. Formaldehyde and HC1 are products.
It is clear that in presence of N02, 03 or hydrocarbon, vinyl chloride is
more photochemically reactive, than in their absence. In experiments
where the hydrocarbon fuel was added, NO disappearance was especially
rapid, confirming that the hydrocarbon supplies significant amounts of
reactive species such as radicals. We might have expected that hydro-
carbon would also protect vinyl chloride by itself competing for the
available reactive species. This evidently did not occur in our systems
to any great extent. Possibly the additional radicals made available by
the presence of the hydrocarbon counterbalanced this effect.
In all runs containing N02, formaldehyde was a product. Approximately
equal amounts of HC1 were formed simultaneously. Gay et al.1-78 ' have
also observed HC1 from photooxidation of vinyl chloride. In their
work, about 50% more HC1 than formaldehyde was produced. They also report
formyl chloride, which decomposes rapidly to CO and HC1. Sanhueza and
Heicklen^99 J. observed HC1 formation from both Cl-atom and 0(3P)-
initiated oxidation of vinyl chloride. They found no evidence for
formyl chloride formation in the former case, but observed it, as well
as CO, HC1 and formic acid, in the latter case. Apparently HC1 is both
a primary and a secondary product.
The major product observed in our experiments was ozone. The peak 0^
values correlated strongly with the initial levels of N02 present. When
292
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296
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hydrocarbon was present, initial ozone production was particularly rapid,
peaking in about 2 hours, as compared with 3-5 hours for runs where hydro-
carbon was absent. This is as expected, as 03 build-up is partly due to
the oxidation of NO to N02 by organic radicals, followed by N02 photo-
lysis. Ozone attack on vinyl chloride is certainly a major factor in the
removal of V.C. in our experiments, as may be seen by the "dark" reaction
(Figure 119), Possibly an intermediate cyclic ozonide is formed, which
can break up into either formyl chloride, CO and H20 or formaldehyde, C02
and HC1. The low levels of HC1 observed, however, imply that other halo-
genated reaction products, not measured in this study must also be formed.
Vinyl chloride would appear to be quite reactive photochemically in
presence of N02 and hydrocarbons, and would presumably be fully photo-
decomposed in the troposphere, though near emission sources it might per-
sist under adverse meteorological conditions.
7.9 Dichloromethane
Simulated tropospheric irradiation for up to 16 hours of dichloromethane
in various mixtures in nitrogen, ultra-zero air, N02 and hydrocarbon fuel
resulted in no significant decrease in dichloromethane concentration. This
compound is evidently quite stable under these conditions. This finding
agrees with those of others, such as Billing et al.f^^), McConnell et al.
(196) and Wilson and DoyleC197).
7.10 PCB's
The following laboratory investigation was designed: (a) to simulate some
important physical and chemical parameters characteristic of both clean
air and polluted air and (b) to study the effect of those parameters on
the rates of disappearance and products formed of a selected chlorinated
biphenyl.
Commercially prepared mixtures of PCB's as mentioned previously are com-
plex mixtures of between 50 and 100 isomers. The use of these mixtures
in photochemical studies poses a nightmare in analyses and interpretation
of product formation. Therefore, we deemed it advantageous to study a
selected isomer. It was estimated that the results of such a study would
yield mechanistically interpretable products.
Ortho-chlorobiphenyl (OCB) was chosen since: (a) it is present in commer-
cial PCB formulations, (b) it is volatile and (c) it exhibits an ultra-
violet absorption similar to other PCB isomers (Figure 3). From a prac-
tical standpoint, the isomer could be obtained from commercial sources in
a very pure form (Analabs, Inc., North Haven, Connecticut).
7.10.1 Results and Discussion -
All experiments were performed in 250 liter FEP Teflon bags, the details
of which have been described (Section 4,1.1). All experiments employed
297
-------
50% relative humidity conditions.
Stability of OCB in the Absence of Ultraviolet Radiation - As indicated
by Figure 120 less than 12% of OCB is lost from the gas phase over a 24
hour period. The 250 liter sample of 1.7 ppm OCB declined to 1.5 ppm in
the absence of ultraviolet radiation. This was found to represent an
unacceptable loss in the light of the experimental design.
Effect of Ultraviolet Wavelength Intensity on Rate of OCB Destruction -
Since OCB absorbs ultraviolet radiation coincident with incident tropo-
spheric radiation, the effect of variation in intensity of those wave-
lengths on OCB was studied in an "inert" atmosphere. Irradiations were
carried out in nitrogen at 50% relative humidity. This was considered
as an inert photochemical environment since neither N2 or H20 absorb
energy in the range utilized. In addition, both require energies greater
than could be supplied by any energy transfer processes to rupture their
bondsC 42 )_ Traces of oxygen present (primarily permeation through Teflon
film) will weakly absorb some long wavelengths (in the visible range) but
represent little potential to effect chemical change in OCB. The bond
energies of N2, H20, and 02 are 226, 119 and 117 Real/mole respectively
which are higher than the highest energy found in the simulated tropo-
spheric ultraviolet radiation utilized G~99 Real/mole). However, the
carbon-chlorine bond, which possesses a strength of 58-97 Real/mole depen-
ding on the carbon's other bonds, does have the potential to rupture under
these conditions. The other bonds in the biphenyl structure are of energy
beyond the tropospheric cut off.
The first experiment utilized radiation which roughly represents the
quality of the tropospheric ultraviolet spectrum. It is difficult to
determine the absolute intensity of this range without elaborate and
expensive instrumentation.
The result of these runs are displayed graphically in Figures 121 and 122
respectively as concentration and log concentration OCB as a function of
irradiation time. A 59% loss of the starting concentration was observed.
The second experiment involved changing the spectral distribution of the
incident radiation. By increasing the lamp distribution from 6 to 24 sun-
lamps, the intensity of the 310 nanometer band is effectively quandrupled
since the blacklamps do not remit in this range.
The results of three runs are displayed graphically in Figures 123 and
124 respectively as concentration and log concentration OCB as a function
of irradiation time. A 97% loss of the starting concentration was ob-
served.
If it is assumed that OCB absorbs radiation in the 310 nanometer range
resulting in chemical change,
1) OCB + hv > OCB*
298
-------
FIGURE 120
^i-iE^YL STABILITY IN 250 LITER TEFLON EAOS
IN TliE ABSENCE OF ULTRAVIOLET RADIATION.
£
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2.0
1.5
1.0
0.5
0.0
0 200 1:00 ?00 800 1000 1200
299
-------
FIGURE 121
GRTHO-CHIOPOBIPHZNYL IS NITROGEN IRRADIATED WITH
SIMULATED TROPOSPKERIC RADIATION (LINEAR PLOT).
2.0
Pu
Pu
>H
§
g
i
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C
0.0
200
600 800
(VI:JUTES)
1000
1200
300
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FIGURE 122
ORTHO-CHLOR03IPHENYL II* NITROGEN IRRADIATED WITH SIMULATED
TROPOSPKERIC RADIATION CLOG CONCENTRATION VS. TTMT?.) _
10.0
1.0
2;
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M
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RUN 1
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PUN 2
PUS 3
600 800 1000 1200
TPTE (''INUTES)
301
-------
FIGURE 125
OPTHO-CHLOR03IPHENYL IN NITROGEN IRRADIATED WITH FOURFOLD
310 NANOMETER INTENSITY SOURCE (LINEAR PLOT).
2.0
0.0
0 200 LOO 600 800 1000 1200
1T*E (MINUTES)
-r -V
302
-------
FIGURE 124
OR7HO-CHLOROBIPHENYL IN NITROGEN IRRADIATED WITH FOURFOLD
310 :;ANO:*ETEP INTENSITY SOURCE (LOG CONCENTRATION vs. TIME).
10.0
g
w
§
g
o
c
o
s
0.01
200 HOO 600 600 1000
TIME (MINUTES)
1200
so:
-------
the rate of disappearance of OCB is expressed as la the rate of absorption
of OCB. la is equal to the product of the molar extinction coefficient e,
the integrated light intensity I., the pathlength L, and the concentration
of OCB. The rate of disappearance of OCB is hence, dependent on the con-
centration of excited OCB*.
If it is assumed that the rate of absorption assuming thermal deactivation
and emission processes are small or constant, then the rate of disappear-
ance of OCB is :
2) d(OCB) * I = el (L(OCB)
~dt~
if we rearrange and integrate to the linear form, the effect of I. can be
evaluated.
3) OCBt
') (OCB)
4)
where t0 = 0.
The rate of 2) is expressed as d the quantum yield of disappearance.
5) log (OCB) = -tI0L
-2303 t + log (OCB)Q
Evaluation of Figures 124 and 122 in terms of 5) for the rate of disappear-
ance eI0L illustrates a rate of 0.948 x 10"^ sec'-'- for the tropospherically
balanced irradiation versus a rate of 4.48 x 10"^ sec~l for the increased
310 nanometer band. Thus a four-fold increase in the 310 nanometer inten-
sity produces a four-fold increase in rate of OCB disappearance.
Effect of Reduced Oxygen - The effect of oxygen in the photodegradation
of OCB was explored by irradiating OCB in air and nitrogen. The nitro-
gen irradiation was viewed as a reduced oxygen experiment since the
initial nitrogen can have up to 20 ppm oxygen and oxygen from room air
permeates through the bag material at a substantial rate. The concen-
tration and log concentration of OCB as a function of irradiation time
are plotted in Figures 125 and 126.
The data indicate a 75% decay of the OCB in air as opposed to a 59% delay
in nitrogen. Again assuming first order kinetics, the semilog plots
yield rates of 2.01 x 10'*' sec'l in air vs. 0.95 x 10"^ sec'l in nitrogen.
The results support the hypothesis that OCB is degrading through direct
304
-------
FIGURE 125
, ORTHO-CHLOROBIPHENYL IH AIR IRRADIATED WITH SIMULATED
TPOPOSPL'EPIC RADIATION (LINEAR PLOT).
£
FU
Pk
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FQ
O
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0.0
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600 800 1000
Tl'ffi (^INUTES)
1200
305
-------
FIGURE 126
ORTHO-CHLOROBIPKEHYL IH AIR IRRADIATED WITH SIMULATED
rROPOSPHERlC RADIATION (LOG CONCENTRATION VS. TIME).
10.0
1.0
M
O
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O
O
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eO.l
0,01
200
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60Q 800 1000 1200
TIME (MINUTES)
306
-------
photodissociation. The slower rate of decay in nitrogen is expected since
the unreactive nitrogen functions as a collisional deactivating body for
the excited OCB*. Collisional deactivation returns the electronically
excited OCB* to its undissociated ground state before photodissociation
can occur.
Oxygen also can function as collisional energy absorber. However, the
photodissociated OCB would probably react readily with molecular oxygen.
Thus, in air the biphenyl radical reacts with oxygen to form a peroxybi-
phenyl radical. In the nitrogen environment the lack of reactive oxygen
also makes more likely the recombination of the biphenyl radical and
atomic chlorines thus producing the observed slower rate of decay.
Chlorinated Products - On the assumption that atomic chlorine is generated
upon the rupture of the weak carbon-chlorine bond, it is believed that the
photodissociated chlorine produces hydrogen chloride under tropospheric
conditions. Wofsy and McElroy in a tropospheric/stratospheric model cal-
culated that kinetic and thermodynamic considerations favor the formation
of HC1 in the troposphere<"198) _ Chlorine dioxide (C102), which is probably
formed readily from the reaction of free chlorine and the abundant mole-
cular oxygen, readily photodissociates under tropospheric conditions.
Other species less likely formed considering the unlikely reaction of two
chlorine atoms are chlorine monoxide (C120) and free chlorine (Cl2). These
would also be readily photodissociated under tropospheric conditions(178,
199). Hydrogen chloride, however, will be readily formed in the presence
of molecular or atomic hydrogen. Hydrogen chloride will not dissociate
under the influence of tropospheric radiation and is stable in the chemical
dynamics of troposphere(198). In the real troposphere HC1 probably reacts
with ammonia to form ammonium chloride. However, in the purified air used,
it was assumed that ammonia was not present (<10 ppb based on NOX measure-
ments) and therefore the chlorine was analyzed as soluble chloride ion
(Cl~) along with hydrogen ion (H+).
Figures 127, 128 and 129 illustrate the corporation of H+ and Cl~ concen-
trations (expressed as micrograms M^) as a function of irradiation time.
Perhaps more revealing is the production of H+ and Cl" normalized against
the OCB lost. Assuming that all OCB lost forms equimolar concentrations
of HC1 then a plot of the data should show a constant ratio (product
observed to product calculated from OCB loss) of 1.0. Figures 130, 131
and 132 illustrate the results for the chlorine yields. Figures 133, 134
and 135 similarly demonstrate the corresponding hydrogen ion yield.
Although a constant 1.0 ratio was not obtained, a few observations are in
order. The scatter observed in the earlier part of the irradiations is
in general much greater than the final data points. This is not unexpected
since the precision of the potentiometric method at low ion concentrations
is poor. Additionally, changes in OCB concentration at high concentrations
are more difficult to detect due to the non-linearity of the electron cap-
ture detector.
307
-------
02
8
O
M
FIGURE 127
SOLUBLE CHLORIDE AND HYDROGEN ION PRODUCTION IN THE
NITROGEN/FOURFOLD 310 NANOMETER INTENSITY IRRADIATION (LOG
CONCENTRATION VS. TIME).
'10-
10
10 —
O RUN 1: CHLORIDE ION
A RUN 2: CHLORIDE ION
Q RUN 3: CHLORIDE ION
O ?U!i I; HYI^OGEN ION
A PUN 2: HYDROGEN ION
PUN 3: HYDROGEN ION
200
600
POO
1000
1200
TIME ("INUTES)
308
-------
FIGURE 128
SOLUBLE CHLORIDE AND HYDROGEN ION PRODUCTION IN THE
NITROGEN/SIMUL/TED TROPOSPHERIC IRRADIATION (LOG CONCENTRATION
VS. TIME).
oo
s
CO
K
O
M
105
10
10-
O RUN 1: CHLORIDE ION
A RUN 2: CHLORIDE ION
D RUN 3: CHLORIDE ION
• RUN 1: HYDROGEN ION
A RUN 2: HYDROGEN ION
B RUN 3: HYDROGEN ION
0 200 1)00 600 800 1000 1200
TIME (MINUTES)
309
-------
FIGURE 129
SOLUBLE CHLORIDE AND HYDROGEN ION PRODUCTION IN THE AIR/
SIMULATED TROPOSPHZRIC IRRADIATION (LOG CONCENTRATION VS. TIME).
O RUN 1: CHLORIDE ION
RUN 2: CHLORIDE ION
U RUN 3: CHLORIDE ION
o RUN i: HYiPOGi::; ION
A RUN 2: HYDPOGSN ION
a RUN 3: HYDP01EN ION
200
1200
310
-------
FIGURE 150
CHLORIDE lONrORTHO-CHLOROBIPHEHYL LOSS VS. TIME IN THE
3ITPO-3KJ/FOURFOLD 310 UANQ'-ETER INTENSITY IRRADIATION.
2.0
co
CO
o
PH
M
pq
o
PC
o
i-3
33
O
O
cc
o
CO
>
o
M
I
K
e
o
M
(X
Q
t-^
W
O
1.5
1.0
„• 0.5
0.0
O RUN 1
A RUN 2
D HUN 3
1
200 100
800 1000
TIME (MIliUTES)
600
9
1200
311
-------
FIGURE 151
CHLORIDE lOiliOPTHO-CHLOROBIPHENYL LOSS VS. TDffi IN THE
UITB03M/SIMULATED TROPOSPHERIC IRRADIATION.
2.0
CQ
CO
PL,
M
pq
o
H
§
C.O
>
o
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K
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1.5
1.0
0.5'
0.0
D
RUN 1
RUN 2
RUN 3
200
Hoo
600
Soo
1000
1200
312
-------
FIGURE 152
CHLORIDE ION:ORTHO-CHLOR03IPHEKYL LOSS VS. TIME IN THE AIR/
SI14ULATED 730POSPHERIC IRRADIATION.
to
s
PH
s
A
as
E3
q
I
§
o
o
M
M
K
s
t^
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O RUN 1
A RUN 2
D RUN 3
200 ^400 600
TI'«E '
1000
1200
313
-------
FIGURE 155
SOLUBLE EYDROGEH ION FOR'ttTIOI?:ORTHO-CELOPOBIPKENYL LOSS
VS. TIME IJ THE NITROGEN /FOURFOLD 310 NANOMETER INTENSITY IRRADIATION.
CO
to
a.
M
CQ
O
K
O
O
K
O
CO
§
M
13
g
g
a
c
O
O
co
2.0
1.5
1.0
0.5
0.0
O- RUN 1
A HUN 2
D • RUN 3
_1
200
600
800
1000 1200
TIME (MINUTES)
314
-------
FIGURE 134
SOLUBLE HYDROSEN ION :ORTHO-CHLOROBIPHENYL LOSS VS. TIME IN
NITROGEN/SriULATED TROPOSPEERIC IRRADIATION.
2.0
CO
M
cq
o
K
I
o
o
•
CO
o
M
PL,
O
M
n
w
I-P
cq
O
CO
1.5
1.0
0.5
0.0
RUN 1
RUN 2
RUN 3
200 1^00 600 800 1000 1200
TIME (MINUTES)
315
-------
FIGURE 155
SOLUBLE HYDROGEN ION:ORTHO-CHLOROBIPHENYL LOSS VS. TIME IN
THE AIR/SIMULATED TROPOSPHERIC IRRADIATION.
u.o
CO
w
(X,
M
pq
o
K
W
O
I
o
EH
g
to
I
o
o
M
O
LO
3.0
2.0
1.0
0.0
O RUN 1
A RUN 2
D RUN 3
200 HOO 600 800 1000
TIME( MINUTES)
1200
316
-------
The final data points, however, are in good agreement for H+ and Cl~. Due
to the poor reproducibility of the earlier data points, however, it is
difficult to discern any distinct trends between experiments. In general,
OCB loss yields a quantitative conversion (50-100%) to HC1.
Irradiations of OCB in the presence of nitrogen dioxide were made but are
not included here because instrument problems rendered OCB analysis poor.
However, it should be noted that phosgene, COC12, was measured at ppb
levels after about 17 hours of irradiation. Phosgene, although not
measured in the experiments presented here, is expected to be produced in
air irradiations without N02.
Phosgene probably accounts for a very small (less than 1%) fraction of the
chlorinated product yield. It is expected that it is formed by the reaction
of atomic chlorine with atmospheric carbon monoxide.
6) Cl + CO > COC1
T) COC1 + Cl -—> COC12
It is not surprising that phosgene is produced in such small yields since
its formation required the collision of CIO and Cl which are both present
at low concentrations. This contrasts with the formation of HC1 which
requires the collision of one chlorine atom with one available hydrogen.
The available hydrogen could be from molecular hydrogen present at 0.5 or
from hydrogen on other PCB molecules.
Oxygenated Products - Since a review of the literature revealed that the
photodissociation of aromatic halides in an oxygen environment generally
produces phenol, aldehydes, and polymers analyses for phenolics and alde-
hydes were performed 0-'^ ) _
Although a detailed aldehydes and phenolics analysis by gas chromatography
was desired, this proved impractical. Phenolics are difficult to separate
by gas unless first derivatized. Higher molecular weight aldehydes like-
wise are difficult to separate without decomposition and the low molecular
weight aldehydes suffer low sensitivity with flame ionization detectors.
Analyses for phenolics and aldehydes were thus performed by wet chemical
methods. Unfortunately, these techniques are non-specific and in the
case of the aldehydes methods, insensitive.
In the irradiations made in nitrogen at the fourfold 310 nanometer inten-
sity, an average of 1.06 ppm aldehydes (at HCHO) remained at the end of
the 24 hour irradiation. The phenolics peaked at 3 ppb in the first 300
minutes and were undetectable at the end of the irradiation (see Figure
136).
In the nitrogen irradiations made with the tropospherically balanced lamp
array (1/4 the previous 310 nanometer intensity), aldehydes averaged
0.45 ppm (as HCHO) at the end of the irradiations. The phenolics again
peaked at around 3 ppb but did not reach its maximum until 700 minutes
(see Figure 137).
317
-------
FIGURE 136
PKIL'JOLICS FORMATION IN THE NITROGEN/FOURFOLD 310 NMO.'-ETER
INTENSITY IRRADIATION.
u.o
CO
o
o
a
PL,
pq
3.0
800 1000
1200
318
-------
rUHB 137
PEENOLICS jOKiATION III THE NITROGEN /SIMULATED TROPOSPHERIC
IRRADIATION.
lt.0
3.0
CO
8 2.0
1.0
0.0 ,
O
A RUN 2
200
J 1
600 800 1000 1200
319
-------
In the air irradiations made with the same lamp balance, aldehydes
averaged 0.37 ppm (as HCHO) at the end of the irradiations. The phenolics
peaked around 200 minutes at about 1.5 ppb (see Figure 138).
It should also be noted that at the end of all irradiations, a character-
istic aldehydic odor was evident. Although this was reminiscent of the
odor of low molecular weight aldehydes, spot checks by gas chromato-
graphy with a flame ionization detector failed to reveal any C2 to €5
aliphatic aldehydes.
The results of the aldehyde production should be viewed in light of the
photochemistry of aldehydes. The CHO group absorbs and dissociates the
a hydrogen in the 310 nanometer band.
Thus, it is surprising that at the fourfold 310 nanometer intensity,
there is more aldehydes remaining than is present at the same nitrogen
irradiation under the balanced lamp array. Since the aldehydes would be
expected to be higher with less than 310 nanometer radiation, the
efficiency of the production of aldehydes is more efficient than its
destruction process.
A comparison of the irradiations conducted in air vs nitrogen reveals
lower aldehydes remaining at the end of the air irradiation. This is
probably a result of further oxidation of aldehydes in the oxygen rich
air.
Figures 136 and 137 demonstrate the phenolics profile observed upon
quadrupling the 310 nanometer band. As can be seen, the phenolics peak
much earlier at the higher 310 band in nitrogen. The production of
phenolics is accelerated by the shorter wavelengths and probably its
destruction rate since it disappears shortly after its maximum is
attained. The presence of oxygen (Figure 138) also increases the rate
of destruction.
Light Scattering Aerosol - Light scattering aerosol measurements on the
air and nitrogen irradiations produced surprising results. Figures 139
and 140 exhibit the formation of aerosol expressed as the scattering
coefficient due to particles in the range of 0.5 to 5 microns (molecular
scattering is small). In terms of scattering coefficient in both irradi-
ations, b scat increased from 0.23 to 100. Converted empirically to mass
loading, this represents an increase from less than 10 micrograms/M to
3-5000 micrograms/M^.
An exploratory experiment conducted at low relative humidity (less than
10%) also showed the same high aerosol formation. Since hydrogen chloride
has a high vapor pressure at room temperature it is unlikely that it would
account for the aerosol observed. Oxygenated aromatics such as 0-hydroxy-
biphenyl, however, have low vapor pressures at room temperature. Thus,
it is not likely from vapor pressure considerations that the light scat-
tering aerosol formed is of organic character.
320
-------
FIGURE 138
PKENOLICS FORMATION IN THE AIR/SIMULATED TROPOSPHERIC
IRRADIATION.
t.O
3.0
8 2.0
pq
fc
A RUN 1
D RUN 3
1000
1200
321
-------
FIGUES 159
LIGHT SCATTERING AEROSOL FORMATION IN THE NITROGEN/
SIMULATED TROPCSPHERIC IRRADIATION (LOG B , VS. TIME).
scat
100
"O
H
•p
a
o
O
10
200
600 800
1000
1200
-------
FIGURE 140
LIGHT SCATTtPIZJG AEROSOL FORMATION IN THE AIR/SIMULATED
TSOPOSPIIZP.IC IRRADIATION (LOG B , VS. TIME).
scat
100
10
-p
K
0
B5
0.1
O RUN 1
A RUN 2
_L L __.
200 1(00 600 600 1000 1200
323
-------
The results of this investigation indicate that 2-chlorobiphenyl is
degraded in the troposphere through photochemical processes to mainly
aldehydes and hydrogen chloride with minor yields of phenolics and phos-
gene.
The observed results indicate that the absorption of radiation of wave-
lengths between 290 nanometers to 310 nanometers leads to the degradation
of 2-chlorobiphenyl. The weakest bond, the carbon-chlorine, is probably
cleaved. This is supported by the appearance of hydrogen chloride and
phosgene.
The bulk of the organic fraction appears to be converted to aldehydic and
phenolic compounds. The detection of total aldehydes as collected by the
bisulfite addition, along with the characteristic odor of short chain
aliphatic aldehydes prompted efforts to pinpoint specific aldehydes.
However, spot checks for acetaldehyde, proprionaldehyde, and benzaldehyde
produced negative results. The exact identity of the aldehyde(s) remains
unknown.
Phenolic compounds detected in minor amounts representing less than 1% con-
version appear to be transient. They build up in the early hours of the
irradiation and die off thereafter. The decrease in concentration may be
due to mechanical removal by condensation and subsequent settling. Also,
it may be removed by further reaction i.e. photolysis, oxidation, etc.
From an overall view direct photolysis in the troposphere may be a very
significant sink for the PCBs. Rates of disappearance for biphenyls of
higher chlorine content will probably be faster than 2-chlorobiphenyl
under similar conditions as evidenced by the liquid phase studies. The
conversion of these compounds to aldehydes, phenols, hydrogen chloride,
and phosgene makes removal by rainout more important since they are more
soluble in water than the parent compounds.
324
-------
LIST OF PUBLICATIONS AND THESES
1, "Absolute Determination of Phosgene: Pulsed Flow Coulometry" by
A. Appleby, H.B. Singh, D. Lillian; Anal. Chem., 47_, 860(1974).
2. "Fates and Levels of Ambient Halocarbons" by D. Lillian, H.B. Singh,
A. Appleby, L. Lobban, R. Arnts, R. Gumpert, R. Hague, J. Toomey,
J. Kazazis, M. Antell, D. Hansen, and B. Scott; A.C.S. Symposium
Series, 17_, 152(1975).
3. "Gas Chromatographic Methods for Ambient Halocarbon Measurements"
by H.B. Singh, D. Lillian, A. Appleby, and L. Lobban; Env. Letters,
(accepted for publication).
4. "Atmospheric Formation of Carbon Tetrachloride from Tetrachloro-
ethylene" by H.B. Singh, D. Lillian, A. Appleby, and L. Lobban;
Env. Letters, 1.0, 253(1975).
5. "Atmospheric Fates of Halogenated Compounds" by D. Lillian,
H.B. Singh, A. Appleby, L. Lobban, R. Arnts, R. Gumpert, R. Hague,
J. Toomey, J. Kazazis, M. Antell, D. Hansen, and B. Scott; Env. Sci.
and Techn., 9_, 1042(1975).
6. Comment on "Atmospheric Halocarbons: A Discussion with emphasis on
Chloroform" by Y.L. Yung, M.B. McElroy and S.C. Wofsy; Geophys.
Res. Letters, 3_, 237(1976).
7. "Gas Chromatographic Analysis of Ambient Halogenated Compounds," by
D. Lillian, H.B. Singh, and A. Appleby; J.A.P.C.A., 26^ 141(1976).
8. "Atmospheric Formation of Chloroform from Trichloroethylene" by
A. Appleby, J. Kazazis, D. Lillian, and H.B. Singh; Env. Letters,
(accepted for publication).
Student Research Theses:
1. J. Kazazis, "Atmospheric Fates of Trichloroethylene and Ethylene
Dichloride."
2. R. Gumpert, "An Investigation of the Ambient Behavior and Atmos-
pheric Chemistry of Carbon Tetrachloride, Methyl Iodide and 1:1:1
Trichloroethane,"
3, L. Lobban, "The Atmospheric Fates of C2C14, CF2C12 and CFC13."
4. J, Toomey, "Atmospheric Fates of Vinyl Chloride and Methylene Chloride,"
5, R. Arnts, "Photodegradation of Ortho-chlorobiphenyl under Simulated
Tropospheric Conditions."
325
-------
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TECHNICAL REPORT DATA
(Please read Instructions on the reverse before completing)
1 REPORT NO.
EPA-600/3-76-1U8
4 TITLE AND SUBTITLE
ATMOSPHERIC FREONS AND HALOGENATED COMPOUNDS
3. RECIPIENT'S ACCESSION NO.
5. REPORT DATE
November 1976
6. PERFORMING ORGANIZATION CODE
1 AUTHOR(S)
Alan Appleby
8. PERFORMING ORGANIZATION REPORT NO.
9. PERFORMING ORGANIZATION NAME AND ADDRESS
Rutgers University
Department of Environmental Science
New Brunswick, New Jersey 08903
10. PROGRAM ELEMENT NO.
1AA603 (1AA008)
11 CONTRACT/GRANT NO.
R-800833
12. SPONSORING AGENCY NAME AND ADDRESS
Environmental Sciences Research Laboratory
Office of Research and Development
U.S. Environmental Protection Agency
Research Triangle Park, North Carolina 27711
13. TYPE OF REPORT AND PERIOD COVERED
Final Report
14. SPONSORING AGENCY CODE
EPA-ORD
15. SUPPLEMENTARY NOTES
16. ABSTRACT
Ambient levels of atmospheric Freons, halogenated hydrocarbons, and SF, were
measured at various locations in the U.S.A. Compounds such as CC1.F, CC12F2,
CH -CC1_, and CC1. were ubiquitious and generally measured at sub ppb levels.
Tropospherically reactive compounds such as C Cl and CHCICCI- were frequently
measured; other compounds were measured where a reasonable source was known. A
novel pulsed flow coulometry gas chromatographic analysis along with other requisite
analytical and calibration procedures were developed and used. Laboratory
irradiation simulations established the tropospheric stability of CC1-F, CC1-F ,
CH CC1_, CC1., CC1 FCC1F , the reactivity of the chlorinated ethylenes, and the
J J *T L. £.
stratospheric reactivity of CC1 F, CC1., and CC1 F . Adventitious labelling of
air masses with halogenated compounds was used to demonstrate urban ozone transport
to rural areas.
17.
KEY WORDS AND DOCUMENT ANALYSIS
DESCRIPTORS
Air pollution
Halohydrocarbons
Atmospheric composition
Chemical analysis
b.IDENTIFIERS/OPEN ENDED TERMS
COSATI Field/Group
13B
07C
04A
07D
18 DISTRIBUTION STATEMENT
RELEASE TO PUBLIC
19 SECURITY CLASS (This ReportJ
UNCLASSIFIED
21 NO. OF PAGES
357
20 SECURITY CLASS (This page]
22. PRICE
EPA Form 2220-1 (9-73)
337
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