United States
Environmental Protection
Agency
Research and Development
Municipal Environments Research FF'A C,00 i 80 1 10
Laboratory ^Vugusl 1 980
Cincinnati OH 45268
Oxidation of
Water Supply
Refractory
Species by
Ozone with
Ultraviolet Radiation
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RESEARCH REPORTING SERIES
Research reports of the Otfice of Research and Development, U.S. Environmental
Protection Agency, have been grouped into nine series. These nine broad cate-
gories were established to facilitate further development and application of en-
vironmental technology. Elimination of traditional grouping was consciously
planned to foster technology transfer and a maximum interface in related fields.
The nine series are:
1. Environmental Health Effects Research
2. Environmental Protection Technology
3. Ecological Research
4. Environmental Monitoring
5. Socioeconomic Environmental Studies
6. Scientific and Technical Assessment Reports (STAR)
7. Interagency Energy-Environment Research and Development
8. "Special" Reports
9. Miscellaneous Reports
This report has been assigned to the ENVIRONMENTAL PROTECTION TECH-
NOLOGY series. This series describes research performed to develop and dem-
onstrate instrumentation, equipment, and methodology to repair or prevent en-
vironmental degradation from point and non-point sources of pollution. This work
provides the new or improved technology required for the control and treatment
of pollution sources to meet environmental quality standards.
This document is available to the public through the National Technical Informa-
tion Service, Springfield, Virginia 22161.
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EPA-600/2-80-110
August 1980
OXIDATION OF WATER SUPPLY REFRACTORY SPECIES
BY OZONE WITH ULTRAVIOLET RADIATION
by
William H. Glaze, Gary R. Peyton,
Francis Y. Huang, Jimmie L. Burleson,
Priscilla C. Jones
North Texas State University
Denton, Texas 76203
R-804640
Project Officer
J. Keith Carswell
Drinking Water Research Division
Municipal Environmental Research Laboratory
Cincinnati, Ohio 45268
MUNICIPAL ENVIRONMENTAL RESEARCH LABORATORY
OFFICE OF RESEARCH AND DEVELOPMENT
U.S. ENVIRONMENTAL PROTECTION AGENCY
CINCINNATI, OHIO 45268
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DISCLAIMER
This report has been reviewed by the Municipal Environmen-
tal 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 endorsement or recom-
mendation for use.
XI
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FOREWORD
The U.S. Environmental Protection Agency was created because
of increasing public and government concern about the dangers of
pollution to the health and welfare of the American people.
Noxious air, foul water, and spoiled land are tragic testimonies
to the deterioration of our natural environment. The complexity
of that environment and the interplay of its components require
a concentrated and integrated attack on the problem.
Research and development is that necessary first step in
problem solution; it involves defining the problem, measuring
its impact, and searching for solutions. The Municipal Environ-
mental Research Laboratory develops new and improved technology
and systems to prevent, treat, and manage wastewater and solid
and hazardous waste pollutant discharges from municipal and
community sources, to preserve and treat public drinking water
supplies, and to minimize the adverse economic, social, health,
and aesthetic effects of pollution. This publication is one of
the products of that research and provides a most vital
communications link between the researcher and the user
community.
Reported here are results of a study of one treatment
technique for the removal of potentially toxic organic substances
from water: the ozone/ultraviolet radiation process. While
still experimental, this process shows good promise as a tool
to achieve the goal of providing pure water for the people of
this country.
Francis T. Mayo, Director
Municipal Environmental Research
Laboratory
111
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PREFACE
The purpose of the work described in this report was to
investigate the efficacy of a new water treatment technology,
the combination of ozone with ultraviolet light. Although the
research itself was necessarily of a very technical nature, the
results are of interest to people with a wide variety of back-
grounds. For that reason, this report is arranged in such a way
that the non-technical reader may read a simple description of
the manner in which the experiments were performed, and a
discussion of the results and conclusions, without being
interrupted by chemical, mathematical and engineering details,
which are collected in the respective Appendices. There are
two exceptions to that format. The first is Section 4, Back-
ground. This section contains material essential to the under-
standing of the ozone/UV process and the details contained in
the Appendices, and is necessarily technical in nature. This
section may be skipped by the non-technical reader.
The second section which does not follow the above format
is Section 7, Process Engineering Analysis and Sizing and
Costing of Full-Scale Units, prepared as a final report to North
Texas State University from Houston Research, Incorporated, the
subcontractor who performed the engineering part of the final
phase of this project. Their final report is included in its
entirety as Section 7, and the non-technical reader may wish to
skip certain sections of that report which deal with mathematical
analysis of the data. The references, whether in Appendices or
the main body of the report, refer to the main reference list.
IV
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ABSTRACT
The use of ozone in combination with ultraviolet radiation
has been studied as an advanced treatment process for the
removal of micropollutants and trihalomethane precursors from
drinking water. The model compounds chloroform, bromodichloro-
methane, tetrachloroethylene, and 2,2',4,4',6,6'-hexachlorobi-
phenyl were treated with ozone/UV as well as ozone and UV
individually in both highly purified water and lake water. The
resulting data were analyzed in terms of kinetic rate expressions
which express the dependence of the reaction rate on ozone dose
rate, UV intensity, and substrate concentration. Similar
studies were performed, monitoring the trihalomethane formation
potential of a natural lake water as a function of ozone/UV
treatment time. The data from the kinetic studies were submitted
to a subcontracted consulting engineering firm, Houston Research,
Inc., for design of a full-scale process and the estimation of
treatment costs.
Products resulting from the ozone/UV treatment of hexachlo-
robiphenyl, chloroform, and the natural organics present in
water from Caddo Lake, Texas, were studied. Complex products
were identified by gas chromatography/mass spectroscopy and a
hexachlorobiphenyl.
The ozone/UV process was found to be 4 to 50 times faster
than either ozone or UV alone, depending on the compound and the
matrix in which it was tested. Photolytic ozonation products of
chloroform were found to be carbon dioxide, hydrogen ion, and
chloride ion. Hexachlorobiphenyl degraded by a process where
one ring appeared to degrade almost completely before the other
ring was broken. Loss of organic halogen occurred very
quickly after ring rupture.
The engineering subcontractor found that water treatment
costs for a 1 to 50 MGD plant to remove 90% of 100 ug/L of
chloroform and bromodichloromethane and 20 yg/L of tetrachloro-
ethylene was $0.063 to $0.16 per thousand gallons. Treatment
costs for removal of 90% of the trihalomethane formation
potential was found to be $0.099 to $0.20 per thousand gallons.
v
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This report was submitted in fulfillment of Grant Number
R-804640 by North Texas State University under the sponsorship
of the U. S. Environmental Protection Agency. This report
covers the period from September 1, 1976, to February 29, 1980,
and work was completed as of February 29, 1980.
VI
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CONTENTS
Foreword iii
Preface iv
Abstract v
Figures ix
Tables xv
Abbreviations and Symbols xviii
Acknowledgement xx
1. Introduction 1
2. Conclusions . 4
Kinetics of ozone/UV reactions 4
Products from ozone/UV treatment of
model compounds 5
Details of the ozone/UV process 6
Engineering analysis 6
3. Recommendations 8
4. Background 9
Use of ozone in drinking water treatment . . 9
Chemistry of ozonation 11
Generation of active species in aqueous
solution 20
Applications of photolytic ozonation in
aqueous solution 27
UV photolysis of halogenated compounds in
aqueous solution 34
Kinetics of the decomposition of ozone and
its reaction with organic compounds
in water 38
Models for ozone mass transfer and
simultaneous chemical reaction 42
Kinetic models for photolytic ozonation in
aqueous solution 45
Conclusions from literature survey 49
5, Chemical Kinetics of Ozone/Ultraviolet-Induced
Reactions of Organic Compounds in Water. . . 52
Experimental Procedures 52
Results and Discussion 54
6. Identification of Chemical By-products from
Ozone/Ultraviolet-Induced Reactions of
Organic Compounds in Water 108
Experimental procedures . 108
Results and discussion 108
vn
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7. Process Engineering Analysis and Sizing and
Costing of Full Scale Units .........
Symbols & nomenclature ............
Process engineering analysis basis ......
Engineering analysis of the reaction data . . 126
Reaction rates for mixtures ......... 1 42
Reactor dynamics — mass and energy transfer. 149
Conclusions from the data analysis ...... 159
Sizing method and cases ........... l6^
Treating unit configuration ......... \
Capital and operating costs ...... • • • 162
Engineering conclusions and recommendations .
References
Appendices
A. Destruction curves for first-order substrates. . . . 190
B. Chemical kinetics of ozone/ultraviolet-induced
reactions of organic compounds in water --
experimental details ................ 207
C. Chemical kinetics of ozone/ultraviolet-induced
reactions of organic compounds in water - details
of kinetic analysis for first-order substrates . . . 220
D. Results of kinetic analysis of ozone and ozone/UV
destruction of hexachlorobiphenyl in water ..... 238
E. Identification of chemical by-products from ozone/
ultraviolet- induced reactions of organic compounds
in water ...................... 256
F. Quality control during the kinetic runs ....... 289
Vlll
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FIGURES
Number Page
1 Direct and radical catalyzed reactions of
ozone in aqueous solution 15
2 Proposed major reaction pathways in the
ozonation of phenol 19
3 Proposed degradation mechanism for the ozonation
of 4,4'-dichlorobiphenyl 21
4 The effect of UV light on ozone decomposition. ... 23
5 Absorption spectrum of O-, from Inn and Tanaka. ... 24
6 Absorption coefficients of C>2 and 03 32
7 Proposed mechanism of 2,4-D photodecomposition ... 34
8 Photolysis of haloaromatics 35
9 UV spectra of four isomeric tetrachlorobiphenyls . . 36
10 UV spectra of chlorobiphenyls 37
11 NTSU photochemical reactor 53
12 Pseudo-first-order rate constants for ozone/UV
destruction of tetrachloroethylene 67
13 Pseudo-first-order rate constants for ozone/UV
destruction of tetrachloroethylene, Detail. ... 68
14 Pseudo-first-order rate constants for ozone/UV
destruction of chloroform 69
15 Pseudo-first-order rate constants for ozone/UV
destruction of bromodichloromethane 70
16 Tetrachloroethylene - effect of pH on treatment
time (95% removal from 100 yg/L) ozone only
(no UV) 72
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Number
17 Effect of pH on ozonolysis of chloroform '
18 Effect of pH on photolytic ozonolysis of chloroform . 75
19 Effect of pH on BDM destruction by ozone ^
20 Effect of pH on ozone/UV destruction of BDM 77
21 Effects of dose rate on the direct ozonation of
HCB in lake water 79
22 Effects of dose rate on the photolytic ozonation
of HCB in lake water 80
23 Ozone and ozone/UV destruction of THM precursors -
Ohio River water (2/77) 82
24 Ozone and ozone/UV destruction of THM precursors -
Ohio River water (2/76) 83
25 Ozone and ozone/UV destruction of THM precursors -
humic acid (Ippm) 85
26 Normalized THM formation potential: Ozone
destruction of THM precursors - Caddo Lake
water 86
27 Analysis of THMFP destruction curves - Caddo
Lake water 87
28 Normalized THM formation potential: Ozone
destruction of THM precursors - Caddo
Lake water 88
29 Trihalomethane precursor destruction by ozone/UV:
[UV] = 0.40 W/L 90
30 Trihalomethane precursor destruction by ozone/UV:
[UV] = 0.196 W/L 91
31 Trihalomethane precursor destruction by ozone/UV:
[UV] = 0.096 W/L 92
32 Mass transfer experiments showing "first run"
and subsequent results 103
33 Production of chloride from the photolysis of
HCB in water 112
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Number Page
34 Hexachlorobiphenyl photolysis pathway according
to Ruzo ..................... 113
35 Proposed degradation scheme for photolytic
ozonation of HCB in water ............ 115
36 Formation of oxalic acid during photolytic
ozonation of water ................
37 Chlorine balance during ozone/UV treatment of HCB. . 117
38 Oxyphotolysis Process Unit for Drinking Water. . .
39 Correlation of Capital and Operating Costs (Assuming
50-60% 03 Utilization) .............. 172
A-l Tetrachloroethylene in purified water. UV
intensity = 0.375 W/L .............. 191
A-2 Tetrachloroethylene in purified water. UV
intensity = 0.189 W/L .............. i92
A-3 Tetrachloroethylene in purified water. UV
intensity = 0.091 W/L .............. I93
A-4 Chloroform in purified water. UV intensity =
0.417 W/L . ................... 194
A-5 Chloroform in purified water. UV intensity =
0.184 W/L .................... I95
A-6 Chloroform in purified water. UV intensity =
0.084 W/L .................... 196
A~7 Bromodichloromethane in purified water.
UV intensity = 0.375 W/L ............. 197
A-8 Bromodichloromethane in purified water.
UV intensity = 0.189 W/L .............
A-9 Bromodichloromethane in purified water.
UV intensity = 0.08.4 W/L ............. 199
A-10 Tetrachloroethylene in Lake Lewisville water.
UV intensity = 0.375 W/L ............. 20°
A-ll Tetrachloroethylene in Lake Lewisville water.
UV intensity = 0.189 W/L ............. 2°1
XI
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Number Page
A-12 Tetrachloroethylene in Lake Lewisville water.
UV intensity = 0.084 W/L
A-13 Chloroform in Lake Lewisville water. _
UV intensity = 0.375 W/L
A-14 Chloroform in Lake Lewisville water.
UV intensity = 0.189 W/L
A-15 Bromodichloromethane in Lake Lewisville water.
UV intensity = 0.375 W/L
A-16 Bromodichloromethane in Lake Lewisville water.
UV intensity = 0.189 W/L
B-l Spectral distribution from ultraviolet lamp. ... 208
B-2 Schematic diagram of the photolytic ozonation
system ..................... 209
B-3 Effect of water purity on photolytic ozonolysis
of chloroform ................. 2 *-•
C-l Determination of order and rate constant of
ozonolysis of tetrachloroethylene ....... 223
C-2 Determination of substrate order in 0->/UV term
(tetrachloroethylene) .............
C-3 Determination of UV intensity exponent and rate
constant for 0.,/UV term (tetrachloroethylene) . 226
C-4 Determination of photolysis rate constant for
bromodichloromethane in purified water ..... 228
C-5 Pseudo-first-order rate constants for bromodi-
chloromethane in purified water ........ 229
C-6 Ozonolysis of tetrachloroethylene in lake water. . 232
D-l Ozone concentration profile for ozonation of
HCB in lake water ............... 245
(
D-2 Ozone concentration profile for photolytic
ozonation under 0.25 W/L UV radiation ..... 246
D-3 Effects of dose rate on the direct ozonation of
HCB in lake water ............... 247
Xll
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Number Page
D-4 Initial rates of direct ozonation of HCB in lake
water at various ozone dose rates 248
D-5 Effects of dose rate on the photolytic ozonation
of HCB in lake water 249
D-6 Effect of UV radiation transfer rates on the
photolytic ozonation of HCB in lake water . . . 250
D-7 pH effects on the rate of ozonation of HCB in
purified water 251
D-8 Effect of pH on rates of photolytic ozonation of
HCB in purified water 252
D-9 Photolytic ozonation of HCB in purified water. . . 253
D-10 Comparison of photolytic ozonation in lake water
and purified water. UV intensity = 0.2 W/L. . .
254
D-ll Comparison of photolytic ozonation in lake water
and purified water. UV intensity = 0.1 W/L. . . 255
E-l Reactor for ozone and ozone/UV treatment of HCB. . 257
E-2 Extraction scheme 259
E-3 GC/MS scan of HCB after photolysis under nitrogen. 262
E-4 GC/MS scan of HCB after photolysis under oxygen. . 263
E-5 Mass spectrum of spectrum no. 834 2^4
"? £ &\
E-6 Mass spectrum of spectrum no. 806 „ .
E-7 Mass spectrum of spectrum no. 903
E-8 GC/MS scan of acid fraction from the photolytic
ozonation of HCB 267
T C Q
E-9 FID chromatogram of HCB/0-/UV reaction mixture . . ^yQ
E-10 HCB photolytic ozonation products at various time
intervals 2-9
E-ll Mass spectrum of spectrum no. 590 270
971
E-12 Mass spectrum of spectrum no. 640 ^ 'J-
972
E-13 Mass spectrum of spectrum no. 690 "' *•
xiii
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Number Page
E-14 Mass spectrum of spectrum no. 700 273
E-15 Mass spectrum of spectrum no. 755 274
E-16 Mass spectrum of spectrum no. 807 275
E-17 Mass spectrum of spectrum no. 820 276
E-18 Reconstructed gas chromatogram; Caddo Lake
water, ozonated. neutral fraction 280
E-19 Reconstructed gas chromatogram; Caddo Lake
water, ozone/UV. neutral fraction 281
E-20 Reconstructed gas chromatogram; Caddo Lake
water, control, neutral fraction 282
E-21 Reconstructed gas chromatogram; Caddo Lake
water, ozonated. acidic fraction (diazomethane) 283
E-22 Reconstructed gas chromatogram; Caddo Lake
water, ozone/UV. acidic fraction (diazomethane) 284
E-23 Reconstructed gas chromatogram; Caddo Lake
water, control, acidic fraction (diazomethane). 285
F-l Linearity of chloroform analysis. Broad Range
(0-400 yg/L) 295
F-2 Linearity of chloroform analysis. Narrow Range
(0-40 yg/L) 296
F-3 Linearity of bromodichloromethane analysis .... 297
F-4 Linearity of bromodichloromethane analysis upon
serial dilution in water 298
F-5 Linearity of hexachlorobiphenyl analysis upon
serial dilution in water (error bars shown for
solution A only) 299
xiv
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TABLES
Number Page
1 Effect of pH on Half-life of Ozone in Water 13
2 Summary of Research on Ozone Decomposition in Water . 39
3 Kinetic Form of Pseudo-First-Order Rate Constants,
K, for Tetrachloroethylene, Chloroform, and
Bromodichloromethane 60
4 Reaction Rate Constants, k^, for Individual Terms in
Rate Expression of Tetrachloroethylene,
Chloroform, and Bromodichloromethane 61
5 Contributions to Pseudo-First-Order Rate Constants
by the Various Component Terms 62
6 Contributions to Pseudo-First-Order Rate Constants
by the Various Component Terms - by Percent. ... 63
7 Ratio of Rate Constants in Purified Water to Lake
Lewisville Water 63
8 Elimination Times, x(3), for Tetrachloroethylene by
Ozone and Ozone/UV for Various pH Values 71
9 TOC Destruction in Cross Lake Water by Ozone
and Ozone/UV 94
10 THMFP Destruction in Cross Lake Water by Ozone/UV . . 95
11 Mass Transfer Coefficients Determined for NTSU
Reactor Compared with Literature Values 105
12 Efficiency of Chloroform Destruction by Ozone/UV. . . 107
13 Products Resulting from Ozone/UV Treatment of
Chloroform 109
14 Characteristics of the Lake Water 126
15 Analysis of Data Sets . 127
xv
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Number Page
16 Summary Data Analysis, Tetrachloroethylene
Destruction in Purified Water .......... 13°
17 Summary Data Analysis, TCE Destruction in
Lake Water (A) .................. 132
18 Summary Data Analysis, Chloroform Destruction
in Purified Water ................
19 Summary Data Analysis, Chloroform in Lake Water. . . 136
20 Summary Data Analysis, Bromodichloromethane
Destruction in Lake Water (A) .......... 138
21 Composite Run, CF , BCM and TCE in Lake Water (A) . . 139
22 Summary of Data Analysis, THMP Destruction in
Lake Water (B) .................. 141
23 TOC Disappearance, Cross Lake Water ......... 143
24 TOC Disappearance, Lewisville Lake Water ...... 144
25 Reaction Parameters, Micropollutants in Lake
Water (A) .................... 147
26 Reaction Parameters, THMP in Lake Water (B) ..... 148
27 Comparison of Reactor Dimensions .......... 149
28 Summary of 0_ Mass Transfer Coefficients and
Transfer Numbers ................ 154
29 Chloroform Mass Transfer Data ............ 156
30 Summary of Design Cases Sizings, Series A ...... 163
31 Summary of Design Case Sizings, B-Series ...... 164
32 Summary of Capital Costs, A-Series ......... 167
33 Summary of Operating Costs, A-Series ........ 168
34 Summary of Capital Costs, B-Series ......... 169
35 Summary of Operating Costs, B-Series ........ 170
36 Operating Cost Estimate Sensitivity, A-Series. . . . 173
37 Operating Cost Estimate Sensitivity, B-Series. . . . 174
xv i
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Number Page
D-l Values of r\f K and K for Ozone Profiles with
and without aUV in Lake Lewisville water .... 242
D-2 Initial Rate, R, and Initial Rate Constant, KQ,
for Ozonation of HCB in Lake Lewisville water . . 243
D-3 Rate Constants, ku, for Photolytic Ozonation of
HCB in Lake Lewisville Water 243
D-4 pH Effects on the Rate Constants, ku, and k of HCB
Destruction by Ozonation and Photolytic Ozonation
in Purified Water 244
E-l Major Acidic Products Identified in Photolytic
Ozonation of HCB in Water 277
E-2 Products Identified from the Ozone and Ozone/UV
Treatment of Water from Caddo Lake, Texas . . . . 286
xvi i
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LIST OF ABBREVIATIONS AND SYMBOLS
ABBREVIATIONS
BDM
CF
CTMA
FID
GC/MS
HCB
L *~
LLW
NTSU
PFO
PW
TCE
THM
THMFP
TOC
UV
W
bromodichloromethane
chloroform
2-chloro-3-trichlorophenylmaleic acid
flame ionization detector
gas chromatography/mass spectrometry
hexachlorobiphenyl
liter
Lewisville Lake water
North Texas State University
pseudo-first-order
purified water
tetrachloroethylene
trihalomethane
trihalomethane formation potential
total organic carbon
ultraviolet
watts
SYMBOLS
Upper Case
'B
D
H
I
K
P
S
S
V
concentration of ozone in liquid
ozone dose rate, mg/L min
Henry's Law constant
— UV intensity, W/L
pseudo-first-order rate constant
trihalomethane precursor; pressure
substrate concentration (also C,A)
initial substrate concentration
volume
Lower Case
La
,P
''purge
rate constant
mass transfer coefficient
ozonolysis rate constant, autooxidation rate
constant
purging rate constant
purging rate constant
XVlll
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k — ozone/UV rate constant
t — time
Greek
a — normalized concentration S/S
r\ — yield factor; also ratio of equilibrium value to
Henry's Law value of ozone concentration
X — wavelength
T -- elimination time
Ji — competition value
xix
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ACKNOWLEDGMENT
Section 7 of this report, Process Engineering Analysis and
Sizing and Costing of Full Scale Units, was prepared by
H. William Prengle, Jr., Arthur E. Nail and Dilip S. Joshi of
Houston Research, Inc., Houston, Texas. Cooperation of the
Houston Research, Inc. team in the initial phases of the
experimental portion of this project is also acknowledged.
Support to the environmental chemistry program at North
Texas State University by the NTSU Faculty Research Committee
is also gratefully acknowledged.
xx
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SECTION 1
INTRODUCTION
Although ozone has been used in Europe to disinfect
drinking water since the turn of the century, chlorination has
traditionally been used in the United States. Within the last
few years it has been shown that chlorination of natural water
produces small amounts of compounds which are toxic or carcin-
ogenic. Other substances undoubtedly are produced, the nature
and toxicity of which are entirely unknown. Thus, alternatives
to chlorine are the subject of many contemporary studies.
To date, applications of ozone in drinking water treatment
have been made primarily in the area of disinfection, but with
its great reactivity, ozone is of potential value in the removal
of undesirable compounds which may be difficult to remove by
other methods. Two major drawbacks to the use of ozone to
remove organic compounds are the cost of the treatment and the
tendency of most ozonation reactions to produce refractory
compounds which are resistant to further degradation by ozone.
The combination of ozonation with irradiation by ultraviolet
light has been reported to overcome the latter difficulty by
degrading even refractory organic compounds completely to carbon
dioxide and water- In addition, the ozone/UV process has been
reported to be considerably faster than ozone alone for the
destruction of almost all organic substrates (103-105). This
increased rate of reaction, coupled with the advances in the
efficiency of ozone generation and contacting expected within
the next few years, could make the ozone/UV process an attractive
alternative for drinking water purification.
This project was conceived because of the need to investi-
gate the possibility of using the ozone/UV process in drinking
water treatment. In the original proposal, six model compounds
of importance in drinking water treatment were chosen as sub-
strates to be studied in purified water. In addition, destruc-
tion of trihalomethane precursors in a natural matrix (river or
lake water) was to be studied. The six model compounds were
chloroform, bromodichloromethane, tetrachloroethylene, a
hexachlorobiphenyl, benzo-a-pyrene, and commercial grade "humic
acid." Kinetic runs were to be performed in the twenty-liter
stainless steel reactor located at Houston Research, Inc.
(Houston, Texas), and samples shipped to North Texas State
University to be analyzed. It soon became apparent that the
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original plan was not feasible, not only because of the volume
of sample handling, breakage, etc., but also because of the
number of runs which would be required to characterize the
extremely complex ozone/UV system. For this reason, a quartz
photochemical reactor, designed along the lines of the sparged,
stirred-tank, semibatch reactor used at Houston Research, was
built at North Texas State University, and was used for all of
the kinetic results presented in this report.
Further modification to the original project plan was
introduced when kinetic runs were attempted using benzo-a-pyrene.
This compound is subject to adsorption onto many materials,
photo-oxidation in room atmosphere and light, and has other
properties which make it an unsuitable compound on which to
perform precise kinetic studies. Furthermore, it was found to
disappear after just a few seconds of ozonation, so that a
treatment process as drastic as ozone/UV was not needed to
remove it from solution. Benzo-a-pyrene was accordingly dropped
from the model compound list.
"Humic acid" was dropped from the list of model compounds
when it was learned from other investigators that the commer-
cially available "humic acid" was not representative of aquatic
humic material. Instead, greater efforts were concentrated on
the study of a natural matrix, water from Cross Lake, Louisiana.
It was realized during the course of the project that an impor-
tant factor related to destruction of trihalomethane precursors
(THMFP) in natural water was the rate of destruction of total
organic carbon (TOC), because of the competition of the "TOG"
matrix with substrate for the active ozone species. Accordingly
the project plan was modified to include the study of TOC
destruction in the same natural matrix for which THMFP was
investigated.
It further became apparent as the project progressed that
the effect of the TOC matrix would also be important in the
destruction of model compounds. The project plan was then
modified in the third and final year to include a complete rerun
of all the purified water kinetics in a natural matrix, water
from Lake Lewisville, Texas. This extensive addition to the
work load necessitated the de-emphasis of product studies, and
consequently only by-products from chloroform, hexachlorobi-
phenyl, and a natural matrix, Caddo Lake, Texas, were
investigated.
The main purpose of this project was to investigate the
efficacy of the ozone/UV process at a laboratory scale by
obtaining kinetic data of sufficient quality that rate equations
could be written which would predict how quickly a given impurity
could be removed by the ozone/UV process, using a given ozone
dose rate and UV intensity. This information could then be
scaled up and optimized in an engineering calculation so that
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the total treatment cost of a proposed process could be projected.
Another important goal of the project was to gain a better
understanding of the ozone/UV process so- that current technology
could be improved to make this potentially powerful treatment
technique more cost effective. This gain in understanding comes
from careful analysis of the data collected from kinetic and
product studies. Finally, as a potential drinking water treat-
ment technique, it is important to investigate the products
which are formed upon treatment of organic compounds with ozone/
UV, in order to avoid problems of the type discovered with
chlorination in the last few years.
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SECTION 2
CONCLUSIONS
The conclusions presented in this report are divided into
four sections:
A) Kinetics of Ozone/UV Reactions
B) Products from Ozone/UV Treatment of Model Compounds
C) Details of the Ozone/UV Process
D) Engineering Analysis
Section C is important since many general conclusions have been
inferred from the kinetic and product work, and the character-
ization of the reactor. These general conclusions are
important in the future consideration of the ozone/UV process
as well as in shaping further developmental work.
KINETICS OF OZONE/UV REACTIONS
1) Ozone/UV treatment was found to be four to fifty times
faster than ozone or UV alone, depending on compound
and matrix.
2) Of the three contributing reactions considered to occur
during an ozone/UV experiment (ozonolysis, photolysis,
and ozone/UV), the ozone/UV term contributes 65-98% of
the reaction rate. In other words, the ozone/UV
reaction is 2-50 times faster than the other important
reaction, photolysis, while ozone appears to contribute
negligibly.
3) All reactions are much faster in purified water than in
lake water, but the compound which photolyzes least
(CF) shows a greater difference between the two
matrices, probably because UV transmission is effected
less than active specie propagation.
4) UV photolysis alone may prove to be an effective
treatment tool, being somewhat specific for some
compounds and not subject to serious interference in
lake water.
-------
5) All three compounds (chloroform, bromodichloromethane,
and tetrachloroethylene) are destroyed more quickly by
ozone alone at high pH than at neutral pH, but the
ozone/UV reactions are all slower at high pH. The
upper limit of the TCE reaction rate with ozone as pH
increases approaches the lower limit of the ozone/UV
reaction.
6) The ozone and ozone/UV reactions of 2,2',4,4',6,6'-
hexachlorobiphenyl are kinetically complex and slow
down significantly by the time substrate is half
destroyed (high kinetic order in substrate).
7) THMFP (Trihalomethane Formation Potential) is seen to
be destroyed more quickly than TOC.
8) Ozone/UV destroys THMFP and TOC faster than does ozone
alone, for a given ozone dose.
9) Ozone/UV appears able to entirely destroy the THMFP of
a natural water. Ozone does not.
10) Precursors in addition to those originally present may
be produced during the ozone/UV process but are later
destroyed as the treatment continues.
11) The rate at which the THMFP of a particular lake water
is destroyed and formed by ozone/UV is very sensitive
to changes in the matrix which may occur over a short
time, either upon storage in the laboratory or in its
natural source.
PRODUCTS FROM OZONE/UV TREATMENT OF MODEL COMPOUNDS
1) Ozone/UV treatment of chloroform results in the
production of the innocuous products carbon dioxide,
hydrogen ion, and chloride ion according to the ratio
03/UV
CHC13 *- C02 + 3H + 3C1~
2) At the high substrate concentration (1.9 g/L) used in
the chloroform product studies, 10.3 chloroform
molecules were destroyed for every ozone molecule
photolyzed. This is indicative of a chain reaction.
3) 2,2',4,4',6,6'-Hexachlorobiphenyl is degraded by ozone/
UV according to a process where one of the two aromatic
rings is ruptured and appears to degrade almost to
completion before the other ring is attacked. One- and
two-carbon fragments are successively split off,
-------
presumably in the form of oxalic acid and carbon
dioxide. One strength of the ozone/UV process over
ozonolysis appears to be the ease of decarboxylation.
Once a ring is attacked, chlorine is eliminated almost
immediately so that the products are presumably more
biodegradable than the parent compound.
4) Ozone treatment of lake water results in many more
chromatographable products than does ozone/UV treatment
of the same water, indicating greater destruction by
ozone/UV.
DETAILS OF THE OZONE/UV PROCESS
1) The purity of the water used to perform ozone mass
transfer experiments can be extremely important when an
iodimetric method is used for ozone determination in
the liquid. Water never before exposed to ozone or UV
gave consistently higher values of the mass transfer
coefficient than the same water in successive runs
after ozone/UV treatment.
2) In the light of 1) above, the NTSU reactor compares
reasonably well with engineering reactors reported
in the literature.
3) Classical ozone mass transfer considerations appear to
be inappropriate for the ozone/UV system, since effi-
ciency calculations indicate that all photolyzed ozone
contributes active species equally, whether photolyzed
in the liquid or in the gas.
4) The autodecomposition of ozone was found to be higher
in the quartz NTSU reactor than that reported by
other investigators (103,145),.
5) On a stoichiometric basis, ozone utilization in the NTSU
reactor for substrate destruction at low concentrations
(100 ug/L) is low. The reasons for this low efficiency
are not known, but definitely warrant study, as a
sizable increase in efficiency would yield a very
powerful treatment tool.
ENGINEERING ANALYSIS
1) The ozone/UV photo-oxidation process can be used
effectively to reduce micropollutants and trihalo-
methane precursors to the desired limits. Simultan-
eously, TOC will be reduced and disinfection will be
accomplished.
-------
2) The cost estimates indicate that daily operating costs,
including amortization of capital equipment and
interest on borrowed capital, are in the range of $0.06
to $0.15 per 1000 gals of water treated, for 10-50 MGD
units using ozone generated from oxygen.
3) A treating cost of $0.06 to $0.15/1000 gals for
combined micropollutant, trihalomethane and TOC removal,
plus disinfection appears reasonable.
4) If higher ozone utilization values can be achieved,
the capital and operating cost could be reduced
significantly; the latter by as much as 30-40%.
5) For raw water streams containing TOC > 10 mg/L, the
possibility of coupling the process with another pre-
TOC removal process to minimize total treating cost
may be desirable. For high TOC streams, each situation
should be examined individually.
-------
SECTION 3
RECOMMENDATIONS
The ozone/UV process shows promise as an advanced treatment
process for the removal of micropollutants and trihalomethane
precursors from water, and should be studied further. It is
effective for removal of compounds such as halogenated organic
compounds, for which other removal methods may be unsatisfactory,
and would be of particular value in water re-use applications.
The mechanism and active species of the process should be
studied in order to enhance the efficiency above that which was
found in this report. Once the identities of the species present
are known, information from the chemical literature may be of
great help in optimizing the process. The conclusions should be
checked at the pilot scale and various methods of contacting
and reaction investigated in order to further optimize the
ozone/UV process, as it appears that a sizable increase in
efficiency may be possible.
-------
SECTION 4
BACKGROUND
USE OF OZONE IN DRINKING WATER TREATMENT
Oxidants are applied in water treatment for several dif-
ferent purposes; namely, for the reduction of taste and odor,
for disinfection of bacteria and viruses, for the reduction of
the total organic content, and for the oxidation of micropol-
lutants. It has been established that the use of oxidants does
not completely remove many organics from treated water but
merely changes their characteristics by partial oxidation or
substitution mechanisms. Five chemical oxidants are commonly
used in water treatment, chlorine, by far the most commonly used
in the USA, ozone, chlorine dioxide, permanganate ion, and
chloramines. Robeck and co-workers (1) reported the results of
experiments with several oxidants and organic compounds. Studies
of the overall impact of the use of chlorine as a disinfectant,
and particularly the by-products of the chlorination process
have increased in number since 1973. Glaze and co-workers (2,3)
and Jolly and co-workers (4,5) showed that numerous chlorinated
organic by-products are found as a result of the chlorination
process, most of which are now recognized to have come from
natural precursors. Early studies of Harlock and Dowlin (6),
and Burtschell and co-workers (7) on the use of chlorine as an
oxidant were generally confined to its effect on taste and ,odor-
causing organics. Morris (8) has reviewed both theoretical and
practical aspects of the chlorination of organic compounds.
Chlorination has been used extensively for water treatment
primarily because chlorine is a powerful disinfecting agent, and
has a relatively long lifetime in drinking water distribution
systems. In 1974, however, Rook (9) and Bellar and Lichtenberg
(10) linked the production of trihalomethanes with the chlori-
nation of naturally occurring humic substances via the haloform
reaction. The pioneering works by Jolly (4,5), Murphy and co-
workers (11), Glaze and co-workers (2,3) have also indicated
that a wide variety of halogenated organic compounds are pro-
duced as a result of the chlorination. Symons and co-workers
(12) extensively discussed the use of this oxidant and the
occurrence of trihalomethanes.
-------
Because of the potential hazard associated with chlorination
by-products-, other methods of water and wastewater treatment are
being examined as alternatives to chlorination. Among the alter-
natives, ozonation has received much attention recently. Ozona-
tion has been used for water disinfection in Europe and at a few
locations in North America. Guinvarch (13) and Lebout (14)
described the experiences of ozonation in France. Rook (15)
reviewed the practice for Europe as a whole. In general, the
European emphasis is on ozone as the primary disinfection method
with the use of a low level of chlorine to provide a residual
disinfection capability in the distribution system.
Although ozonation has been used in treatment in Europe
since the turn of the century, economics have usually favored
other methods in North America. However, recent improvements in
ozone generator technology (16) have lowered the cost consider-
ably relative to chlorination, and many workers are looking with
renewed interest at the advantages of ozonation.
Ozonation deserves consideration as an alternative to
chlorination because it has many of the properties required for
an oxidant used in water treatment for removal or organic com-
pounds. It is known to be a more powerful disinfectant than
chlorine. Katzenelson and co-workers (17) demonstrated that
with the same concentration, 99% kill for poliovirus 1 is reached
in less than 10 seconds with ozone. Chlorine requires about 100
seconds to achieve the same effect. Ozone is a powerful
oxidizing agent with twice the chemical oxidation potential of
hypochlorite ion. As a result, more complete oxidation can be
expected from ozonation than from chlorination. It is effective
at relatively low doses and short contact times. Ozone in water
readily decomposes to oxygen leaving no by-products that need to
be removed.
McNabney and Wynn (18) demonstrated at the Blue Plains Pilot
Plant the reduction of chemical oxygen demand (COD) by ozonation.
Numerous studies also revealed the ability of ozone to reduce
total organic carbon (TOC) and biochemical oxygen demand (BOD)
of wastewater. Jiejers (19) showed the reduction of color of
ground and lake water by ozonation. A review by Rosen (20)
gives an excellent overview of the application of ozonation for
the removal of a variety of pollutants including potential
carcinogens from drinking water.
Ozonation also is applied for tertiary treatment of muni-
cipal wastewater in combination with other treatment methods,
such as mechanical filtration (21) or ultrasonics and catalysts
(22). Elia and co-workers (23) used ozonation as a polishing
process for the pre-treatment of reverse osmosis influents.
Another study by Foulds (24) combines ozonation with froth flo-
tation for disinfection and pollutant removal. An encouraging
work by Guirguis and co-workers (25) demonstrated improved
10
-------
performance of activated carbon by pre-ozonation. In their
pilot plant, ozone pre-treatment rendered certain organic
species more sorbable and biodegradable. The ozonated effluent
contains dissolved oxygen (DO) in sufficient concentrations to
promote the growth of active micro-organisms in the filter unit
and on the carbon surface in the sorption unit. A similar study
is currently under extensive investigation by Glaze (26).
Several inorganic substances can essentially change the
properties of water, such as smell, color and taste. These
include compounds of iron, manganese, arsenic, sulfur, and
several others. Removal of these substances can be effected
significantly by ozonation. Shambaugh and Melnyk (27) found
that ozonation with contact times on the order of 10 minutes will
destroy metal-EDTA complexes. In the exhaustive ozonation of
seawater, Williams and co-workers (28) observed a rapid oxidation
of Br~ to BrO~, with the evidence for the transitory formation
of C1CT. Removal of manganese and iron also has been demon-
strated by Mignot (29) . The treatment of wastewater containing
lead has been developed by Collier (30). Lapidot (31) reported
on the removal of silicones .from a silicone manufacturing
effluent. Ozone was the only treatment technique that was shown
to clean up the silicone waste significantly.
Lawrence (32) found the pre-ozonation reduced the concen-
tration of chloroform precursors so that less chloroform was
detected after ozonation-chlorination than by chlorination alone.
Glaze and co-workers (33) used the same procedures on river
water and reported the reduction of trihalomethane formation
potential (THMFP) by ozonation. In another study, Riley and
co-workers (34) found that ozonation-post-chlorination may lead
to higher levels of trihalomethanes than the conventional
chlorination process. However, attention should be paid to the
fact that pH played an important role in their study. The
enhanced chloroform yields were found only if chlorination was
done at high pH following ozonation. Other similar findings
have been documented in the report of a pilot plant study (12).
Although it was proposed that ozone oxidizes organic materials
to form precursors which may, when subsequently chlorinated,
produce chloroform, the complexity of humic material makes the
chemistry of ozonation-post-chlorination difficult to investi-
gate directly.
CHEMISTRY OF OZONATION
The chemical reactions between organic compounds and ozone
have been studied extensively for many years, but the chemistry
of ozone as an oxidizing agent of organic and inorganic solutes
in water is relatively unknown. Studies of the nature of
compounds formed during ozonation of wastewaters are currently
underway to clarify the mechanism of its action.
11
-------
At room temperature Ozone (03) is a highly unstable,
colorless gas which is 1.5 times more dense than oxygen and which
thermally decomposes to oxygen with critical decomposition
temperature at 270°C (35,36). In water, ozone has an oxidation
potential of 2.07 volts at 20°C, which makes it capable of oxi-
dizing most organic and inorganic compounds. The microwave
spectrum of ozone shows the molecule to be non-paramagnetic with
an obtuse angle of 116° 49' and equal bond length of 1.278 A
(37). On the basis of this structure, in nonaqueous solution or
as a free molecule, ozone is expected to react as an electrophile
or as a 1,3-dipole.
The aqueous solubility of ozone is ten times greater than
that for oxygen. Kinman (38) reported 575 and 775 mg L -1 as the
theoretical solubility at 20 and 10°C under one atmosphere of
pressure and at normal pH. However, since ozone is generated in
the form of a dilute mixture (1-5% W/W) in_air or'oxygen gas,
solubilities in water in excess of 20 mg L are seldom,
attained in water treatment practice.
The decomposition of ozone in aqueous solution is well
known and is the subject of a number of studies. Ozone decom-
poses rapidly in water with a half life of a few minutes. The
decomposition of ozone in water is favored at alkaline pH values.
Table I shows the results obtained by Stumm (39) and Hoigne and
Bader (40). Both the mechanism and kinetics of the dissociation
of ozone in water are uncertain. Kilpatrick and co-workers (41)
and Stumm (42) stated that ozone decomposes to yield oxygen and
a small amount of hydrogen peroxide. Garbenko-Germanov and
Kozolva (43) had investigated the decomposition of ozone in basic
medium using the electron spin resonance technique and also
absorption spectroscopy. They suggested the initial step of the
decomposition probably is the reaction between ozone and
hydroxide ion to form ozonide ion and hydroxyl radical (OH-):
03 + OH' ^ °3~ + °H*
The hydroxyl radical then reacts further. This process would
explain the increased dissociation of ozone with increasing
alkalinity. Recently, Hoigne and Bader (40,44,45,46) have shown
experimentally that as pH increases, the kinetics of ozonation
of organic compounds changes. They found reaction rates for
ozone oxidation reactions at high pH are similar to those
obtained in studies in which the hydroxyl radical is generated
by high energy radiation. Hydroxyl radical may be the impor-
tant active species in ozonation as they concluded.
Peleg (47) reviewed the chemistry of ozone in the treatment
of water and from several individual studies and findings, he
suggested the following steps for ozone decomposition in aqueous
solution:
12
-------
TABLE 1. EFFECT OF pH ON HALF-LIFE OF OZONE IN WATER
PH
4.0
6.0
7.6
8.0
8.5
8.9
9.2
9.7
10.0
10.4
Half-Life (min.)
Stuiran (39) Hoigne and Bader
14.6°C 25 C
• ' 350
' ' 50
41 ' '
' * 33.3
11 * *
7 ' '
4 * '
2 ' '
.33
0.5 * *
13
-------
O + HO _ *- 0_ + 20H
J ^ £t
03 + OH- - ^- 02 + H02'
03 + H02- - *~ 202 + OH"
OH- + OH
OH' + H02 -- ^^ H2° + °2
OH- + OH~ - ^- 0~ + H2O , [6]
0~ + 02 - *- 03~ [7]
H0- + H0- - »- H0 + 0 [8]
In this mechanism, 0 has lower chemical reactivity which
differs markedly from that of OH- in many reactions. The half-
life of ozonide, 03~, at 25°C is only several milliseconds and
the substance is reported to be unreactive to aromatic molecules
as well as to methanol and ethanol (48) . Barr and King (49)
showed the hydroperoxyl radical (HO--) to be almost inert towards
organic substances. It would appear then that the hydroxyl
radical is very likely the most important reactive intermediate
in the decomposition of ozone. It is an even more powerful
oxidizing agent than ozone with an oxidation potential of 2.80
volts at [H ] = 1.0 M. (50). Consequently, one may speculate
that the efficiency of the ozonation process might be enhanced if
means were found to accelerate the decomposition of ozone so that
hydroxyl radicals could be used effectively. Kewes and Davison
(51) experimentally proved that the rate of organic oxidation
increases as conditions are utilized which accelerate the
decomposition of ozone.
Hoigne and Bader (40,44-46) stated that both molecular ozone
and hydroxyl radical, which is generated from the decomposition
of ozone, are involved in the ozonation process. The primary
oxidation initiated by ozone in water can be described by the
reaction sequences shown in Figure 1. The overall effect
initiated by ozonation is a superposition of the "direct
reaction" and the "radical-type reaction." Hoigne and Bader (147)
later developed the kinetics formulation for the elimination of
solute (M) in a batch or plug-type reactor:
-d[M]/dt = -n d[0,]/dt = k[0,][M]
J «J
and from integration:
(
14
-------
-ln([H]e/[M]Q) = n • k • [03] • t
where ,
n = yield factor; or the amount of M eliminated per
mole of ozone used,
k = rate constant
[M] , [M] = initial and final concentration of substrate assuming
CO-,] = mean ozone concentration, during reaction period t
03 stripped
0
added
Figure 1. Direct and radical catalyzed reactions of ozone in
aqueous solution (Hoigne', 1976-7) .
15
-------
Oxidation by molecular ozone is highly selective and is
often rather slow (minutes). However, hydroxyl radicals are very
reactive towards most organic solutes in water. The amount of
OH radicals which are available for oxidation of a solute M
depends on the amount of OH formed and the relative rate with
which they react with M when compared with the rate by which
hydroxyl radicals are consumed by all other solutes:
Rote of OH Consumption
+ M
*"" k ' LM] [OH]-
[OH]-
where 03/A = 0 decomposed during process
n1 = yield for OH formation from O-,,A
k' = 2nd order reaction rate constant for OH
Z(k.[S..]) = rate of OH scavanging by all solutes, S.,
present including ozone and M
Only a fraction (n11) of the OH radicals reacting with solutes
will result in a solute elimination.
The rate of solute oxidations in the presence of competing
scavengers becomes:
- ~^f = n' n" (d o3>A/dt) - k; [M] (Z(k; [S±]) r1 [9]
if k [M]«Z(k.[S.]), then integration of equation [9] yields for
a plug and batch type reactor:
- in [M] /[M] = n'n1' o, . • k.. • (Z(k. [s.]) )~i
eo -J,AJXL 11
This equation and the assumptions involved in its derivation
were successfully tested, since rate constants for OH radical
reactions with hundreds of aqueous solutes are known from
independent measurements in which the radicals are generated by
high energy photolysis of water (40).
16
-------
The direct reaction of ozone with organic substrates in
nonaqueous systems has been well studied (52). Bailey (53)
discussed the organic groupings which can be oxidized by ozone.
Generally, carbon-carbon multiple bonds and aromatic molecules
are very easily attacked by ozone. In aqueous ozonation, although
the mechanism is not clear yet, one would expect molecular ozone,
hydroxyl radicals, oxygen and even water molecules to be involved
in the reaction. Consequently, one would expect final products
to differ in aqueous ozonations as compared to those reported in
nonaqueous solvents. For example, hydroxyl hydroperoxides or
dihydroxy peroxides might be produced in the reaction of ozone
with carbon-carbon double bonds as the well-recognized Criegee
mechanism (54) shows:
°3
Rf* — f*D ^^_ o ^ OD __^^ta_ ^\ f\ f* D _L D ^ -• ^
pv — Wr\ ^ n pL> wH -^ w""U~w ri * n^v/ — ^
R
(CARBONYL OXIDE)
H20
OH
0 = C~ •< H-0-O-C-R
^R R
(HYDROXY HYDROPEROXIDE)
0}
H •
H90 * ,C-OH
R
Aromatic compounds, amines, acetylene, and many other com-
pounds have been subjected to ozonolysis in aqueous solution but
few definitive mechanistic studies have been made. The results
of these studies have been reviewed by Bailey (53), Oehlschaeger
(55), Maggiolo (56) and Rice (57).
Ozonation of natural organic ("humic") material in water
has been the subject of several investigations, but few of these
report on the products of ozonation. Guirguis and co-workers
(25) monitored the relative molecular weight distribution in
wastewater before and after ozonation and found the concen-
tration of UV-absorbing species of relative molecular weight
between 500 and 180 increased as ozone dose rate increased.
17
-------
Apparently, the larger humic molecules were cleaved at several
points in the ozonation process. For identification of the
possible small products, Kinney and co-workers (58,59) ozonized
humic acids prepared from oxidized bituminous coal. They stated
that both carbonic and oxalic acids appeared to be primary
products and accounted for about 65% of the carbon. Sievers and
co-workers (60) found unidentified aldehydes from the ozonation
of biologically-treated secondary effluent from wastewater.
Lawrence (61) ozonized the predominant classes of compounds
found in natural waters: humic and fulvic acids, tannins and
lignins and the products were identified by gas chromatography.
He found humic and fulvic acids gave very similar oxidation
products — namely aliphatic and aromatic carboxylic acids of
molecular weight less than 370. Tannic acid and lignosulfonic
acids also yielded carboxylic acids, but mostly of low molecular
weight. Shevchenko and Taran (62) identified aprocrenic, crenic
and oxalic acids as ozonation products from peat humid acids.
Aromatic compounds are very common as pollutants in water
and wastewater. Two types, phenols and chlorobenzenes, are of
interest. Most aromatic compounds are reactive towards ozone
although the reaction is said to be slower than that of olefins
(63). Substituents which withdraw electrons from the ring, such
as carbonyl, halogens, nitro and sulfonic acid groups, deactivate
the ring towards attack by ozone. Electron-releasing substi-
tuents, such as hydroxyl, alkyl and alkoxyl groups, activate the
ring towards attack by ozone. Phenols which are attacked very
rapidly have been studied extensively because of the recent
interest in the use of ozone to purify wastewater. Gould and
Weber (64) ozonized aqueous phenol solutions and found catechol,
hydroquinone, glyoxal, oxalic acid and glyoxylic acid. Yamamoto
and co-workers (65) carried out the ozonation of a variety of
compounds in water. The major products from the ozonation of
phenol, they stated, were formic acid with minor amounts of
hydroquinone, catechol, muconaldehyde, muconic acid, maleic
aldehyde, glyoxal, glyoxlic acid, oxalic acid, carbon dioxide,
and hydrogen peroxide. The proposed major reaction pathways
are shown in Figure 2.
Shuval and Peleg (66) found about 80% of chlorine on
o-chlorophenol was converted to chloride ion upon ozonation
indicating that covalent carbon-chlorine bonds are broken with
ozone. It was also suggested that the active oxidizing species
attacks the ring at a site or sites other than the chlorine site,
because an induction period was observed during which the con-
centration of o-chlorophenol decreased but no chloride ion was
formed. After all of the o-chlorophenol had disappeared,
chloride ion still was being produced upon continued ozonation.
Gilbert (67) also observed the same phenomenon. Chloride ion was
found only after 40% of the o-chlorophenol had been destroyed.
18
-------
A. Hydroxylation ( ) and Cleavage ( ) of Aromatic Ring
OH
i
I
OH
OH
t
^/COOH
OH
OH
t
OH
OH
OH
I
B. Normal ( ) and Anomalous (-
urated Aldehyde and Acid
-) Ozonolysis of a,g-Unsat-
CHO
c
SS5-
CHO
COOH
COOH
-J
-1
u
r1
>-
^^^ ^^ ^J ^^
^^ n \j
^0 ..
--KM
^COOH
(^* ^1 ^J
^^ fi ^^
(CHO
z^
COOH
i CHO
i
—'CHO
H-^COOH
1 1 i
1 1
11 » HC-O-O
CHO 2 HCOOH
JJ ^HC-0-6 ^ MCOOIII
COOH " "vwv"1
ft
Figure 2.
Proposed major reaction pathways in the ozonation of
phenol in water (65) (Continued-next page).
19
-------
C. Oxidation of Aldehyde and Decarboxylation of Formic Acid
CHO CHO COOH
CHO COOH COOH
HCOOH *- C02+H20
(Figure 2. Continued)
However, 4-chlorophenol produced chloride ion at the start of
ozonation. The difference in the rates of dechlorination might
be due to the different distributions of electron density on the
aromatic rings. Besides chloride ion, Gilbert also detected
formic acid and oxalic acid from the ozonation of 2,4-dichloro-
phenol.
Bauch and co-workers (68) detected chlorotartaric acid,
o-;m-and p-chlorophenols and chloride as the ozonation products
of chlorobenzene. The hydroxylation of the ring by ozone is
interesting since it activates the aromatic ring towards further
degradation. Aliphatic oxidation products similar to those
found from phenol were also found by these workers.
Ozonation of one of the major classes of pollutants,
polychlorinated biphenyls (PCBs), has been studied by Yokoyama
and co-workers (69). An aqueous mixture of PCBs, KC-300, was
ozonated for 24 hours and the products were monitored by gas
chromatography. Major products were chlorobenzoic acids,
glyoxylic acid, oxalic acid, chloride ion, hydrogen peroxide,
carbon dioxide and white precipitates which showed strong
carbonyl group vibrations in the infrared spectrum. A
degradation mechanism of 4,4'-dichlorobiphenyl was proposed
(see Figure 3) . According to this view, ozone attacks the II-
electrons of the phenyl nucleus and causes cleavage of the
nucleus. It is interesting that only one ring was ruptured in
many cases, and that the ring ruptured products are similar to
those of phenol and chlorophenols. No direct evidence to
support the reaction scheme was given. Moreover, the active
specie was proposed to be molecular ozone without any confirm-
atory evidence.
GENERATION OF ACTIVE SPECIES IN AQUEOUS SOLUTION
From the previous review, it is generally recognized that
both the decomposition products of ozone and ozone itself are
involved in the oxidation process. In many cases, the former
20
-------
are more important; therefore, better utilization of ozone may
be accomplished by more efficient generation of decomposition
products which are reactive towards organic substrates.
Adjustment of pH has been proven to be effective in
accelerating both the decomposition of ozone and the ozonation
rates of organic substrates (40,42,47,51,70). Higher reactivity
of ozone under basic conditions is now understood to involve the
formation of hydroxyl radical intermediate which is more reactive
than molecular ozone in general. However, except in few cases,
complete removal of organic compounds at higher pH has not been
obtained.
Cl
C = 0
C = 0
H
C-OH
C = 0
+HCI
H
/H
o=c-c + co,
HO 0
Figure 3. Proposed degradation mechanism for the ozonation of
4,4' - dichlorobiphenyl.
21
-------
The effect of temperature on ozone decomposition has been
studied by Kuo and co-workers (71) . They demonstrated that the
Arrhenius equation could be employed to express the temperature
dependence of the decomposition rate constant of ozone. Hewes
and Davison (51) observed a similar dependence on temperature
and found experimentally that ozonation at elevated temperature
reduced the COD of pretreated secondary municipal wastewater
effluent more effectively. Auguliaro and Rizzuti (72) studied
the temperature enhancement of the ozonation of aqueous phenol
solutions. For actual water treatment processes, however, the
practice of elevating the temperature in order to provide better
oxidation is impractical. Also, the low solubility of ozone at
higher temperatures might adversely affect the rate of organic
oxidation.
Chen and co-workers (73) developed catalytic ozonation in a
steady state reactor using a special ferric oxide catalyst. The
authors stated that the catalyst enabled them to use at least
two oxygen atoms in the ozone molecule for oxidations. They
postulated that in catalytic ozonation ozone molecules are
adsorbed on the surface of the catalyst to form ozone-catalyst
complexes which then would form reactive oxygen species for
oxidation. The removal of COD in aqueous systems by ozone with
glass beads and ozone with catalyst, F&2°3' were compared in
this work. The latter system effectively removed COD at a rate
of about three times that of former system. The better utili-
zation of ozone was proven, but the cost and life time of the
catalyst system was not well-investigated. Moreover, inhibition
of the catalyst by chemical poisons is likely to occur in the
treatment of wastewaters containing complex components.
Ultraviolet (UV) radiation is one of the most effective
methods to accelerate the decomposition of ozone. Ozone absorbs
UV light and decomposes very rapidly in water. In a typical
reactor, ozone transfers from the gas phase to the liquid phase
to reach an equilibrium value. When the UV light is turned on,
the concentration of ozone drops almost instantaneously to a
very low value. (Figure 4).
The decomposition of ozone by UV in the gas phase has been
studied intensively. Ozone absorbs radiation in two regions,
one in the red part of the visible spectrum Uma--600 nm) and
the other in the ultraviolet (X.max-254 nm) . The photochemical
behavior is distinct in the two regions (74) . From the UV-
absorption spectrum by Inn and Tanaka (75) the large absorption
coefficient of ozone at 254 nm" is noted (Figure 5) .
Heidt and Forbes (76,77,78), Norrish and Wayne (79)
studied the UV-photolysis of dry ozone and found the quantum
yield is in excess of two. Addition of water vapor gives even
higher quantum yields. Kistiakowsky (80), Castellano and
Schumacher (81) found the quantum yield for the decomposition of
22
-------
to
u>
QC
UJ
I
UJ
O
N
O
6
or
LU
O
O
O
0
10
20
30
40 50
TIME (min.)
90
Figure 4. The effect of UV light on ozone decomposition,
(Ozone Dose Rate = 7.2 mg/min-L, Flow Rate =1.0 L/min).
-------
a.
o
160
E
o
•5
S 80
u
40
(a)
(b)
I
2000 2400 2800 3000 3100
Wavelength, A
(c)
3200 3300
(d)
E
o
a>
'u
a>
o
o
c
o
o
-------
°3
0*
OH1
HO,
H2°"
0,,
0.
2 OH-
OH- +20,
[11]
[12]
[13]
[14]
where 0* is in the D excited state. In contrast to the O( P)
ground state which has all orbitals occupied and which will tend
to act as a radical, 0(1D) corresponds to an electronic struc-
ture in which a 2p orbital is vacant, so that it is expected to
be a strongly "acidic" reagent (eq. [12]). Fortin and co-workers
83) and a number of other researchers (84,85,86) confirmed the
reaction of O(-'-D) with H20 to be one of the principal mechanisms
for producing hydroxyl radicals.
The photochemistry of ozone in water presumably is somewhat
similar, but the large number of water molecules will quench the
excited ozone molecule and consequently the quantum yield is
expected to be lower in the liquid phase. Ozone decomposition
by photolysis to form hydrogen peroxide in aqueous solution has
been reported by Taube (87) . In mildly acidic solution (dilute
hydrochloric or acetic acid) which inhibits the chain reaction
of 0 and HO~, the reaction
and H2O2,
0^ -
hv
H2°2
[15]
takes place with almost exact stochiometry. The quantum yields
for the disappearance of ozone at A=254 nm, 310 nm and 600 nm
are 0.62, 0.23 and 0.002-0.005 respectively. When light in the
longer wavelength region is absorbed, no hydrogen peroxide is
formed. The above observations are in complete agreement with
predictions made on the basis of the gas-phase photochemistry of
ozone discussed earlier. Reaction [ll] occurs followed by direct
formation of hydrogen peroxide by the addition of O( D) to H20.
Tracer experiments have shown that H202 formed derives approxi-
mately one-half its oxygen from the water, the other half from
the ozone. Taube concluded that in mildly acidic system, a
mechanism with OH- as an intermediate would not yield a linear
build-up of H202/ nor would the stoichiometry of eq. [13] be
followed. Thomas (88) suggested that OH- radicals might dimerize
to form H202 which consequently would react with more OH to form
HO,,- and H20 in the absence of acids.
On the other hand, H~02 itself can be photolyzed to
generate hydroxyl radicals. Baxendale and Wilson (89) showed
at high intensities that
hv
2 OH-
[16]
25
-------
is the primary process of the decomposition. Volman and Chen
(90), Hunt and Taube (91) demonstrated that allyl alcohol, a
radical scavenger, can remove the OH- and further reduce the
quantum yield. They found at A=254 nm, in aqueous solution,
the quantum yield for the decomposition of hydrogen peroxide is
0.94 + .06, while in the presence of allyl alcohol the number
drops to 0.54 + 0.05. Tetrahydroxyhexane was found as a product
which must be formed by the addition of hydroxyl radicals to
allyl alcohol followed by the association of the intermediate.
Confirmation of the formation of hydroxyl radicals from the
photolysis of H202 at X=254 nm was also made by Buxton and
Wilmarth (92).
Hydroxyl radicals can be obtained from the decomposition of
H2°2 by other means; such as the titanium trichloride method of
Dixon and Normann (93); generation by ferrous ion in the famous
Fenton's reagent (94,95), and high energy radiation of aqueous
solutions as by Currie and Dainton (96).
As one may expect, H2O2 ^s caPable of acting as an oxidant
in aqueous solution. Recently, Ogata and co-workers (97) studied
the photo-induced oxidation of diethyleneglycol dimethyl ether
(DEDE) and analogs with aqueous hydrogen peroxide. They stated
that the oxidation is initiated by hydrogen atom abstraction by
hydroxyl radical formed by the photolysis of H209. This is a
model reaction for the photo-oxidative removal of nonionic sur-
factants in wastewater which are difficult to remove biologi-
cally. The absorption coefficient for H202 is larger than that
of DEDE at 250 and 300 nm, so it is assumed the photolysis of
DEDE itself is negligible. The hydroxyl radicals are supposedly
the only active specie to react with DEDE. Ogata and co-workers
reported that destruction of up to 67.2% of substrate could be
reached after an irradiation time of 60 min. for a mixture of
0.937 M H902 and 0.184 M DEDE in aqueous solution at 20 C. The
products round were methanol, ethylene glycol monomethyl ether
(EME), methoxylacetaldehyde, methoxyacetic acid and methyl
formate. These presumably were formed by the abstraction of
hydrogen atoms at various sites followed by bond cleavage. No
complete removal has been reported by this process.
Dore and co-workers (98) used the UV/H202 system for the
oxidation of phenol solutions and compared the results with
ozonation. Solutions were chlorinated after treatment and
chloroform yields determined. It was reported that both systems
gave the same result in that the quantity of chloroform produced
by post-chlorination passes through a maximum value which
depends on the time of pre-treatment. Photolytic oxidation by
UV/H202 thus was proved to have similar effect'on the degradation
of organic compounds as ozonation. One may expect, however,
that the quantum yield of hydroxyl radicals from hydrogen
peroxide might be too low to effectively oxidize organic com-
pounds for practical application. The difficulty in preparation
26
-------
and storage of hydrogen peroxide also discourages the use of
this oxidant in large scale water treatment.
APPLICATIONS OF PHOTOLYTIC OZONATION IN AQUEOUS SOLUTION
The ultraviolet-ozone (03/UV) system has been used for
oxidation of dieldrin in nonaqueous solution by Nagl and Korte
(99), and the oxidation of carbon tetrachloride and chloro-
fluoromethanes in the vapor phase by Jayanty and co-workers (1.00).
More recently, application of the process in aqueous solutions
have been developed by a group at Houston Research Inc. It has
been established that photolytic ozonation rates of dissolved
organic compounds are ten to one thousand fold greater than with
ozonation alone. Garrison and co-workers (101) applied 03/UV
for the destruction of complexed cyanides in water. They
reported that for a fixed concentration of ozone, increasing
the intensity of UV light increased the reaction rate markedly.
Likewise, for a fixed intensity of UV light, increasing the
ozone concentration also increased the reaction rate markedly.
Ozonation alone is very slow compared to any combination of ozone
concentration and UV intensity studied. Fochtman and Huff (102)
reported that wastewater containing small concentrations of TNT
was completely oxidized to carbon dioxide with ozone while being
irradiated simultaneously with ultraviolet light. When ozone
only was applied to three liters of "pink" water which contained
63 mg/L organic carbon, the TOC dropped to 56 mg/L over the
first 60 minutes. No further decrease was observed, even after
an additional 2.5 hours of ozonation. Under the same conditions/
but with UV radiation, the total organic carbon decreased to
17 mg/L in a period of two hours. From the carbon mass balance
studies. Fochtman and Huff found that 85% of the carbon lost
from TNT solution under O3/UV was accounted for as C02- Within
experimental error, they concluded that TNT had been completely
destroyed. No light molecular weight organics were found.
Prengle and co-workers (103) studied the rates of oxidation
of refractory chemical species, ethanol, acetic acid, glycine,
glycerol and palmitic acid, with ozone and ultraviolet light.
A Refractory Index (RFI) parameter was defined to give a
measure of the difficulty of oxidation of a given molecule.
The lower its RFI, the less refractory is the compound to
ozone oxidation. The authors monitored the changes of TOC to
indicate the destruction rate of each substance with 03/UV.
The results were as expected, in that ozone together with UV
light enhanced the destruction significantly. In discussing
the mechanisms for the oxidation processes in the 0~/UV system,
the authors postulated two models of the kinetics or 0^/UV
process. When the amount of UV radiation is small, and the
chemical reactions are rate controlling, the authors visualized
the reaction as being predominantly a three-step sequence
consisting of simultaneous interaction of substrates with ozone
and UV light. However, when the intensity of UV radiation is
27
-------
larger, the reaction becomes mass transfer controlled and
substrates may be photolyzed to form free radicals which then
react with ozone to form C02 and H20 as final products. Appar-
ently the authors applied "classical ozonation" in the model _
and omitted the possibility of ozone being decomposed to provxde
active species which carried the major reaction. Ozone/UV
oxidation of several chlorinated compounds also had been studied
by Prengle's group (104). They examined the oxidation rates of
the following five compounds: pentachlorophenol,_o-dichloro-
benzene, dichlorobutane, chloroform and polychlorinated
biphenyls, Aroclor 1254. In general, none of the above compounds
are oxidized to the desired lower limits of concentration by
ozone treatment alone, but are effectively destroyed by 03/UV.
The original and partial oxidation species reportedly are
completely oxidized to Cl~, Cl2f C02, and H^O. It was also
concluded that UV absorption correlates with reactivity, e.g.,
aromatics absorb mor£ UV and are more reactive than^aliptiatic^
compounds. The fate of chlorine on~these compounds was described
by one of two~^foces¥esldechlorination to form chloride ion
from those compounds with substantial UV absorption, or
transformed to chlorine and then to hypochlorpus^acid. The
latter"case reportedTy~^cciirred wherTTETie compound had sjmalj.__UV
^absorp_tion_ such as chloroform. TSfcTevidence was given to "^verify
^EEeTorination of free chlorine. Later, Prengle and Mauk (105)
carried out photolytic ozonations of four pesticides: Malathion,
Baygon, Vapam and DDT in aqueous solution. All were destroyed,
as well as subsequent partial oxidation products, to yield TOC
values below detectable limit. Postulated mechanisms were
presented which indicate the chemical routes from initial to
final oxidation species. The authors stated that O^/UV photo-
oxidation of dissolved substance (M-species) occurs primarily by
a combination of ozone photolysis and M-photolysis. For low
le ve 1 __UV_ absorber s the oxidation rate is con trpl led_by_thej
|>hot6Tyinrs~~oT~ol:qi^^ so ~s~tated~~tEat the species,
~H."ariH~OH radicais~~formed from photolysis of water, as well as
H and OH , are involved in the photooxidation of substrates^
Neither the intermediate nor the final residues from photolytic
oxidation of any of the organic compounds discussed were
characterized.
The reduction of COD in wastewater by 03/UV has been demon-
strated by Wako and co-workers (106). A wastewater with 40.2
ppm COD and 110 ppm dispersion dye was treated with ozone while
being exposed to UV light at 210-350 nm. After treatment the
water had a COD of 2.8 ppm. When the wastewater was treated
with ozone alone, the treated water had a COD of 37.3 ppm. In
another report, Wako (107) studied the photooxidation of phenol
with ozone and UV radiation. The yield of decomposition prod-
ucts was determined by liquid chromotography. The products
were not identified but the relations between the peak heights
of products and reaction time for each decomposition were
recorded. The possible intermediates of ozonation of phenol,
28
-------
such as catechol, hydroquinone, resorcinol and maleic acid were
also subjected to C>3/UV at various pH values. As expected,
decomposition of phenol was found to be accelerated by increasing
the concentrations of base (potassium hydroxide).
Ozonatbons with and without UV irradiation of 2-propanol,
acetic acid, acetone and oxalic acid were conducted by Kuo and
co-workers (108) . Ozonation and C>3/UV products from these model
compounds were identical. Upon ozonation, 2-propanol was con-
verted to acetone which in turn was oxidized to acetic and
oxalic acids. Trace amounts of formaldehyde and formic acid
were detected in the ozonated acetone' solution. Further
ozonation of acetic acid resulted in the formation of glyoxylic
acid that was oxidized readily to oxalic acid which in turn was
oxidized directly to carbon dioxide. The destruction rate of
each compound was observed to be faster during Oo/UV process
than with ozone alone. The results of Kuo and co-workers suggest
a very interesting aspect of the O3/UV processes for the photo-
oxidation of organic compounds. It has been recognized that
saturated carbonyl compounds undergo photo-induced decarboxy-
lation in the gas-phase. The process was first observed by
R.G.W. Norrish and is known as the Norrish Type I or a-cleavage
process (109). The solution-phase reaction is less common due
to more rapid vibrational relaxation by collision with solvent
molecules, but still many reports of the reaction occurring in
solution phase have been made:
0
n
R1CR2 + hv *•- R^' + R2C = 0 [ 17]
• *
or R,C =0 + R2'
The free radicals formed are believed to react with solvent or
other radicals. In other cases, once the carbonyl groups are
excited they commonly undergo intramolecular hydrogen abstrac-
tion from the side chain by the photo-excited carbonyl group with
the formation of a 1,4-biradical. This species can be cleaved
to an olefin and an enol. These reactions are now referred to
as Norrish Type II reactions:
OH
CHR2
29
-------
Acetone, then, may be oxidized by the Norrish Type I
process followed by reaction with OH- from the decomposition of
ozone under UV radiation as follows:
0
II
hv
CH_-C=0 +
3 i l "
| OH• |OH •
CH3COOH CH3OH
[19]
Acetic acid may be oxidized in the following manner (110) :
OH- CHO OH'
CH3COOH
CH2OH
COOH
hv
CHO
I
COOH
hv
COOH
I
COOH
[20]
OH
hv
In accordance with the postulated pathway, all the inter-
mediates except CH-OHCOOH were identified during 0.,/UV of
acetone by McCarthy (111).
The fate of methanol under O^/UV was shown by the experi-
mental data of Chian (112). By successive abstraction of
hydrogen, addition of hydroxyl radicals or oxygen, and
rearrangement of unstable intermediates, methanol may be
oxidized through formaldehyde to formic acid. Further oxidation
of acids would result in the production of carbon dioxide and
water:
OH-
OH '
OH •
CH3OH
HCHO
HCOOH
C0
[21]
hv
hv
hv
Complete removal of the substance thus occurs once the
decarboxylation occurs.
Sine tana and Bulusu (113) studied methods of removal of RDX,
hexahydro-1,3,5-trinitro-l,3,5-triazine, from wastewater of
munitions manufacturing and processing plants. They found that
aqueous RDX solutions undergo rapid photochemical degradation
with 254 nm radiation leading to low molecular weight .gaseous
products. Ozone was found to degrade RDX much less efficiently
than UV alone, and the combined use of ozone and UV exhibited a
marked synergistic effect in that the combined degradation rate
is greater than the sum of the two individual degradation rates.
30
-------
Total destruction of substrate was observed, with C00, NO, NO,
CO, HCN, CH20 and H2O found as the products.
In addition to their work on ozonation for destruction of
halomethane precursors from humic acid and river water, Glaze
and co-workers (33) also applied 03/UV in the same system. It
was found that both ozone and 03/UV effectively reduce THM
formation potential. O3/UV was found to be more effective than
ozone alone.
Leitis and co-workers (115) studied the nature of oxygen-
ating species by determining the ratio of hydroxylation products
formed by ozonation and 03/UV. They attempted to show that
hydroxyl radical was not the active specie in ozone and ozone/
UV treatment by comparing the relative proportion of ortho-,
meta-, and para-nitrophenols which were found as products using
ozone, ozone/UV, UV alone, hydrogen peroxide, Fenton's reagent,
NaOCl/H^O-j, and H202/UV both with air and nitrogen present.
Although very little data was given in the paper, it appears
that the authors were basing their conclusions on products which,
although purportedly representing the first of a multi-stage
process, accounted for only about 1% of the substrate which had
been destroyed up to that point. The relative amounts of those
products remaining would of course be related to their relative
rates in subsequent processes, as well as to the manner and
amount in which active specie was generated, the initial con-
centration of substrate, and the length of the treatment time.
Since the ozonation was performed at approximately neutral pH,
hydroxyl radical is not expected to play an important part in
their experiments using ozone without UV, and the results are
therefore not in contradiction with the conclusions of Hoigne
and Bader (44-46).
The above authors found that ozonation and 0.,/UV of
nitrobenzene, selected as a primary model compound, gave similar
ratios of all major peaks in the chromatograms of the reaction
mixtures from ozone and ozone/UV treatment. The major stable
intermediates were oxalic acid, formic acid and glyoxal. It
was stated that the reactivities of some of the model compounds
during both processes were similar indicating that the same
reactive species is involved in the reactions. However, the
reactivities of the intermediates, oxalic and formic acids, were
greatly enhanced by 03/UV over ozonation. The authors concluded
that the first steps of the ozone and ozone/UV degradation are
probably not free radical ift nature, but gave no direct evidence
other than the fact that the reactivities of nitrobenzene,
benzole acid, and anisole with ozone/UV were inconsistent with
electrophilic aromatic substitution.
In the treatment of water by ozonation and combined ozone/
ultraviolet radiation, ozone at 1-5% by weight in air or oxygen
is normally applied at a certain rate into the reactor. Under
UV radiation, there remains the question as to whether the
31
-------
interaction of oxygen and ultraviolet radiation contributes to
the destruction of substrates. In order to examine this ques-
tion, two groups, Fochtman and Huff (102) and Prengle and co-
workers (105) conducted experiments in which oxygen was used to
replace the ozone-oxygen mixture for photolytic oxidation.
Fochtman and Huff reported that very little effect could be
noted for the destruction of TNT. However, Prengle and co-
workers found that oxygen with low pressure mercury radiation led
to some degree of destruction of reactant species, such as
pesticides. They stated that the effect was a result of a
combination of mechanisms which included the production of trace
amounts of ozone, activated oxygen, and activated reactant
molecules.
The absorption spectrum of oxygen obtained by McNesby and
Okabe (116), Figure 6, shows that it strongly absorbs radiation
SCHUMANN-RUNGE
BANDS
1200
1600 2000
X, A
2400
Figure 6_. Absorption coefficients of 02 and Oo
32
-------
of wavelengths below 176 nm. Although oxygen weakly absorbs
radiation within the forbidden Herzberg band, originating near
254 nm (117),
02(3E~g) + hv »- 02(3£+u) ^- 2 0(3P) [22]
only ground state oxygen O ( P) atoms, not excited molecular
oxygen can be formed to act as active oxidative species (eq.[22]).
Oxygen must absorb radiation with wavelength below 176 nm to
gain enough energy to be promoted to the 02(3Z~u) state which
may subsequently dissociate to form 0(1D):
02(3Z~g) + hv ^- 02(3Z~u) ^-0(1D) + 0(3p) [23]
Compared with the facile decomposition of ozone at 254. nm, the
dissociation of oxygen is far less important in the 0-,/UV process.
On the other hand, molecular oxygen can undoubtedly take part
in the reaction when substances form radicals by other pathways.
Galen and Smith (118) have studied the photodecomposition of
phenol in a tubular-flow reactor in which phenol was dissociated
initially to form the phenoxy radical. The authors propose that
this radical then was attacked by molecular oxygen to form
peroxy radicals which ultimately yield organic acids and other
products.
Jacob, Balakrishnan and Reddy (148) proposed mechanisms
which explained the production of phenol and hydroxymucondialde-
hyde from the reaction of benzene with hydroxyl radicals in
which molecular oxygen played a very important role. Correspond-
ingly, Mathews ar*d Sangster (149) showed a threefold decrease in
carbon dioxide yield from the hydroxyl radical-involved decar-
boxylation of benzoic acid when the aqueous solution was
deoxygenated. In fact, for reactions of hydroxyl radical in
dilute aqueous solution, the oxygen molecule is probably a very
important contributor, since it provides the mechanism by which
an intermediate can rid itself of the unpaired electron; for
example (148) ,
+ H02
33
-------
UV PHOTOLYSIS OF HALOGENATED COMPOUNDS IN AQUEOUS SOLUTION
The nonoxidative reactions of organic compounds in aqueous
solution caused by UV-irradiation would also be expected to occur.
The principles of photochemical reaction in aqueous solution have
been reviewed by Owen (119). Water provides a transparent and
sometimes reactive medium in which chemical reactions are effec-
tively energized by UV. Ultraviolet light has been shown to
exert drastic changes in many pesticides under laboratory con-
ditions. It is well known that, in the presence of hydroxide
ions or other nucleophiles, photonucleophilic displacement
reactions result in hydrolysis and the conversion of aromatic
halides to phenols and eventually to insoluble polymers. Crosby
and Tutass (120) studied the photodecomposition of 2,4-dichloro-
phenoxyacetic acid (2,4-D) at 254 nm. The decomposition products
found were 2,4-dichlorophenol, 4-chlorocatechol, 2-hydroxyl-4-
chlorophenoxyacetic acid, 1,2,4-benzentriol, and polymeric
humic acid." The proposed mechanism is shown in Figure 7.
C'\
OCH2COOH
OCH2COOH
OH
Figure 7. Proposed mechanism of 2,4-D photodecomposition.
Apparently, from the excited state, the carbon-halogen bond
undergoes fission giving rise to aryl and halogen radicals. The
radicals then abstract hydrogen from the medium or dimerize.
Prior to bond fission an alternative reaction between the excited
state and a nucleophilic species may occur giving the appro-
priate substitution at the carbon-halogen bond (121).
34
-------
Figure 8 shows the nonoxidative photolysis of haloaromatics
Cl
Figure 8. Photolysis of haloaromatics
Polychlorinated biphenyls are known to be stable towards
oxidation, hydrolysis and other chemical reaction might occur
in the environment. However, the photochemically-induced
degradation of PCBs has been reported by a number of groups,
such as Ruzo and co-workers (122-124), Safe and Hutzinger (125,
126) and Hustert and Korte (127). Most of the photochemical
experiments were conducted in organic solvents such as hexane
and methanol. Dechlorination is the common reaction observed.
The isomeric effects play an important role in the rate of
dechlorination of PCBs. The spectra of these compounds show a
wide range of absorptions which are related to the individual
isomeric structure. MacNeil and co-workers (128) examined UV
spectra of twenty-nine chlorobiphenyls. It is noted that with
increasing chloro substitution in which there are less than two
chlorine groups ortho to the ring-ring bond the ^ma^ values for
the K bands are shifted to longer wavelengths. Introduction of
two or more chlorine atoms ortho to the ph-ph bond results in
a decrease in the intensity of K band (see Figure 9 and 10).
35
-------
Figure 9.
195 220 245 270 295 320 nm
UV spectra of four isomeric tetrachlorobiphenyls,
36
-------
Cl Cl Cl Cl
cone. = x80
cone. = x80
cone. = x80
i I I 1 L
200 260 320
Figure 10. UV spectra of chiorobiphenyls.
37
-------
Since biphenyl itself has a planar excited state geometry,
under radiation the coplanarity between the ring would increase
conjugation and stabilize the excited state molecules. However,
the ortho chlorine substituents sterically hinder the coplanarity
between the phenyl rings, and these groups are specifically
cleaved in preference to substituents at the met and para posi-
tions. An increase in ortho substituents decreases the lifetime
and increases the reactivity of the excited state.
It was found that oxygen retards the photochemical degrada-
tion rate of PCBs in aerated solution. Ruzo and co-workers (123)
also observed that 1,3-cyclohexadiene quenched the excited state
of 2,2',4,4'-tetrachlorobiphenyl in degassed methanol solution.
Both results indicate the excited state is a triplet. The authors
concluded that PCBs should be excited from ground state to
excited singlet and to excited reactive triplet state from which
radicals are formed.
2,2',4,4',6,6'-Hexachlorobiphenyl (HCB) had been irradiated
at 310 nm in hexane by Safe and Hutzinger (125). They found some
products formed by loss of chlorine, rearrangement and conden-
sation; however, the products were not well-characterized.
As noted, most photochemical reactions of PCBs have been
conducted in nonaqueous solution. In an attempt to demonstrate
the photochemical reaction of PCBs in "natural condition" which
is close to that in the environment. Hutzinger and co-workers
(126) irradiated PCBs in aqueous suspensions and films covered
with water. They identified hydroxychlorobiphenyls and "barboxyl"
type compounds, which showed hydroxy and carbonyl stretching,
besides dechlorinated PCBs which were identified by gas chroma-
tography, mass spectrometry and infrared spectrometry.
KINETICS OF THE DECOMPOSITION OF OZONE AND ITS REACTION WITH
ORGANIC COMPOUNDS IN WATER
The rapid decomposition process of ozone in aqueous
solution has been studied by many investigators, as reviewed in a
previous section. These investigations have included a number of
studies on the kinetics of ozone decomposition; however, the
findings are conflicting. In particular, there is a disagreement
over the kinetic order of reaction.
Autodecomposition of ozone in water, as it is often called,
has been reported to follow a variety of kinetic orders which are
apparently pH dependent. A summary of the range of variables
covered by various investigators and their conclusions concerning
the reaction order related to ozone is shown in-Table 2.
There is presently no general agreement on the mechanism of
autodecomposition nor has there been any adequate explanation of
the variations in the kinetic results obtained by various
38
-------
TABLE 2. SUMMARY OF RESEARCH ON OZONE DECOMPOSITION IN WATER
u>
vo
Investigators
Alder and Hill
Chang
Czapski and co-workers
Hewes and Davison
Kilpatrick and co-workers
Kuo and co-workers
Merkulova and co-workers
Rankas and co-workers
Cizzufi and Marrucci
Rogozhkin
Rothmund and Burgstaliker
Sennewald
Shambaugh and Melnyk
Stumm
Weiss
(129)
(130)
(70)
(51)
(41)
(71)
(131)
(132)
(133)
(134)
(135)
(136)
(137)
(42)
(138)
ph Range
1-2.8
9
10-13
2-4, 6, 8
0-6.8, 13
2.2-11.0
0.22-1.9
5.4-8.5
8.5-13.5
9.6-11.9
2-4
5.3-8
9.0
7.6-10.4
2-8
Temperature range , C
0-27
25
25
30-60,10-50,10-20
25
15-35
5-40
5-25
18-27
25
0
0
25
1.2-19.8
0
Reaction order with
respect to ozone
1
1
1
2, 3/2-2,1
3/2,2
3/2
1,2
3/2
1
1
2
2
1
1
3/2
-------
investigators. Several investigators include pH dependent terms
in the rate expression such as Stumm (42) .
-°-75 C24]
= k [OH-]-[0]
and Weiss (138)
L3] = kn [OH"] [O,] + k, [OH~]0'5 [0^]3/2 [25]
dt 13^ -j
The mechanism proposed by Peleg which combines proposals by
Stumm and others is shown in equations [1] through [8] .
The kinetics of the ozonation of some substances in water
and wastewater have been studied. Khandelwal and co-workers (139)
reported that cyanide removal appeared to follow one-third order
with respect to cyanide concentration:
-dd[CN"] = k [ON']1/3 [26]
where k the reaction rate constant contains substantially con-
stant ozone concentration in the liquid. Kirk and co-workers
(140) proposed the following expressions based on their experi-
ments of the removal of COD by ozonation. The removal or
organics followed either the pseudo-first-order kinetics as
shown in equation (27)
= [COD] [27]
_
or second-order kinetics as in (28):
[COD] [28]
In [27] and [28] k.. is the pseudo-first-order rate constant,
[COD] is the concentration of COD at time t, k~ is the second-
order rate constant, [0^] is the dissolved ozone concentration
at time t.
\
Chang (130) as well as Kuo and Wen (141) studied the
kinetics of the ozonation of methanol and its degradation
products in aqueous solution. They obtained the following
40
-------
expression for the ozonation of methanol:
- d [CH3OH]
—g£ = k2[03]P[CH3OH] [29]
where p = 3/2 to 1/2 as pH varied from 7.2 to 11.0. Their data
also showed that the ozonation of formaldehyde and formic acid,
degradation products of methanol, are also first-order with
respect to the substrates, however, the order with respect to
ozone concentration varies. For formaldehyde
- d[HCHO]
at = k[03]q[HCHO] [30]
where q = 1/2 as reported by Kuo and Wen (141) and'one by Chang
(130). For formic acid
= k[0]r[HCOOH] [31]
3
where r = 1.2 to 0.74 as pH increases from 1.7 to 11.0. With
substrate at an initial concentration, the degradation should be
proportional to the substrate concentration and to the quantity
of ozone used to degrade it.
Eisenhauer (142) defined the "ozone dose rate" as follows:
[0] F
C32]
where O- = ozone concentration in the feed gas (moles/L)
F = ozone gas flow rate (L/min)
P = initial phenol concentration (mole/L) , and
o
V = reaction volume (L) .
*
Thus, the degradation rate of phenol could be expressed as
-_|_[P3 = k.R.[P] [33]
where [P] = concentration of phenol at time t. Since phenol [P]
is oxidized to catechol (C) which in turn is degraded to other
41
-------
oxidation products (D), Eisenhauer used the kinetics of consec-
utive reactions of the type:
kl k2 [34]
The rate expression for concentration of catechol at time t may
be derived from this expression; yielding 35 :
[35]
where P , R and F are defined in equations [32] and [33] .
Gould aSd Weber (64) in another study of the ozonation of phenol
also found that the degradation of phenol follows the expression
[33] proposed by Eisenhauer. They stated that the removal of
COD may be expressed by the similar relationship
[COD] = [COD] e~k " R ' t [36]
o
MODELS FOR OZONE MASS TRANSFER AND SIMULTANEOUS CHEMICAL REACTION
Since ozone is only slightly soluble in aqueous solution,
the rate of absorption of ozone is of main concern in consid-
ering the ozone injection process. Eisenhauer (142) as well as
Gould and Weber (64) had demonstrated the dependency of ozonation
rate on the ozone dose rate. Mass transfer, therefore, plays an
important role in the semi-batch type reactor system. The
relationship between the dose rate and dissolved ozone concen-
tration is the subject of the study of mass transfer. Numerous
studies have been done on models for ozone mass transfer (137,
141, 143-145). The mathematical model which can describe the
mass transfer process of ozone from air or oxygen into the water
is analogous to heat transfer or conduction of electric charge
through resisting medium. Film theory is commonly applied for
the mass transfer model, according to which, the ozone reaction
mechanism for treating water can be visualized as follows for
the case where the ozone-substrate reaction is very fast (145) :
a) Ozone diffuses through the gas film to the gas-liquid
interface.
b) Ozone transfers across the interface into the liquid
film.
c) The liquid reactant diffuses from the bulk liquid into
the liquid film.
42
-------
d) Ozone oxidizes the liquid reactant to form a product.
e) The product can continue to react or diffuse into the
bulk liquid,
Using film theory, Shambaugh and Melynk (137) developed the
following model to describe the concentration of ozone, A°, in
the liquid as a function of time.
| mo [l-e~PRh] GHe ,. -PRh, ..-
A° = _V - . l-e- kl + VP- (1'e
_ -,
J
1 VP
where
_^ rkL kGa' , [38]
Gh LkT + krH J
L G e
G = inert gas flow rate, g mole/sec.
V = volume of reactor, cm
m = molar fraction of ozone in the gas feed
P = reaction operating pressure, atm.
h = distance from sintered glass plate to surface of liquid
in reactor., cm.
k, = first-order rate constant for ozone decomposition, sec .
kT = liquid-film mass transfer coefficient, cm/sec.
L
2
k_ = gas-phase mass transfer coefficient, g mole/cm -sec-atm,
G
a' = interfacial area per unit volume of liquid, cm
He = Henry's law constant, atm-cm /g-mole.
Prom equation [37], the steady state concentration of ozone can
be expressed as
A°
ss
| . m .
= V o
k + GHe
r, -PRh-,
L n-0-pRhi
43
-------
A°ss is less than the equilibrium ozone concentration predicted
by Henry's law due to rapid ozone decomposition, particularly
under alkaline conditions (pH = 9 and above).
In an attempt to obtain k a, the overall liquid-phase mass
transfer coefficient, Yocum (145) derived the following expres-
sion also using film theory:
kTa =
LI
n - 3 c.
-] -
J V
[40]
with
-1 XB
Cl = £ ln Cl - xf
[41]
and
n = i +
X
[42]
Be
where X = mole fraction of ozone in liquid
X
B
*
B
theoretical mole fraction of ozone in liquid at
equilibrium with feed.gas as predicted by Henry's
law
X
B = mole fraction of ozone in liquid at equilibrium
k = ozone decomposition reaction rate constant, Ib.
mole/ft min.
L = moles of liquid, Ib moles
V = volume of liquid in reactor, ft .
6 = PtL/HG
Pt = total pressure, atm.
H = Henry's law constant, atm.
G = gas molar flow rate, Ib mole/min.
This coefficient is also a function of turbine power input and
superficial gas velocity.
44
-------
Equations [40] - [42] were derived in part from the
equation which describes the change of ozone concentration in
the liquid. Writing slightly differently from reference 145,
dCB
d!T= Kla(CBe - CB) ~ VB ' C43]
where C_ = bulk concentration of ozone in liquid,
CB = concentration of ozone in liquid, which would be
in Henry's Law equilibrium with the gas,
and other symbols are as before. Mass transfer coefficients
were calculated from the measurement of ozone adsorption into
distilled water as a function of time.
When other substances are present which could also react to
remove ozone from solution,
OZONE REMOVAL RATE. = k.SaC^ = -R. , [44]
1 1 1 ri 1
dCR
diT - Kla (CBe - CB> ' VB - f, ' [45]
When there is only one term in ER_ and it is first order in CB
then -. i i
ER = k C_ [46]
i i a B
and
n = (1 + ko + ka) = C* [47]
As will be seen later, this is the case for ozone photolysis.
KINETIC MODELS. FOR PHOTOLYTIC OZONATION IN AQUEOUS SOLUTION
Very little work has been done on the kinetics of ozonatLon
under UV irradiation. According to Wako (107) , the reaction
rate of the 03/UV oxidation of phenol can be expressed as
= k'[P]
°'5 = "P°'5 [48]
45
-------
where k1 = 0.63 exp (-3. 18xl03/T) , k1 '=1. 24xl04exp (-3J.8X103 /T)
P = concentration of phenol in solution, mol/L
03 = the ozone dose rate, mol/L (02) /sec.
T = temperature of solution
Since the decomposition is also reported to be catalyzed
by base (potassium hydroxide) , Wako derived the following
equation for the disappearance of phenol under 03/UV:
[P]KOH=0~[P] = CP]t=c){l-exp(-1.3xl03[KOH][03]t)}. [49]
Prengle and co-workers (105) developed kinetic expressions
for the O3/UV process. The expressions were based on the assump-
tion that reaction mechanisms were different under low and high
intensities of UV radiation. The authors stated that when the
UV intensity is low, then
k k k
A r— 2- - ^- p l »» p 2 *- P_(C00+H_0) [50]
V°3 l V°3 2 hv'°3 3 2 2
and
TOG
= noe ~^16 " (V^l~1)e
where w , w, , w« are the reaction rate moduli containing k , kn ,
k2, and0 ^ * o 1
wl W2 wo W2
^o E (wo-wo) (w2-wQ) ' fi 5 (Wl-w0) (w2-W;L) [52]
However, when UV is larger,
A = hv - *- R.'s [53]
R-'s + 03 - *- C02 + H20 + . . . [54]
and the generation rate of free radicals is
46
-------
_ dR- 'S _ U(j,a I
dt " -
C55]
where kya is the radiation transfer coefficient
expression obtained is
The final
[56]
where 6 = V (Sr~) *vc' fl(t) = 1~et'
Ao
f2(t) =
~wot -
~Wlt -
~w2t,
input
is the critical
When UV input is small, f (t) = 1 and
) = f (t) , and when
UV input is larger, t is small, f 2 (t) = 1 and
) = f 1 (t) .
No investigations have been reported in which there is
sufficient detail of data to confirm this model or to evaluate
the reaction rate moduli, but it would appear that the validity
of the assumption in equation [53] could easily be deduced from
the results of photolysis experiments on the individual
substrate compounds.
Taking a completely different approach from that of
Prengle and co-workers, Lee and co-workers (110) developed a
mathematical model describing the O,/UV removal of total
organic carbon (TOC) from water. The authors assumed that under
UV radiation, the only important active species are hydroxyl
radicals which are generated from the decomposition of ozone
which previously had transferred into the liquid phase.
According to this view, the radicals either react with the
organic substances or react with each other to terminate the
chain reaction. The decomposition rate constant k
-------
a) When the total organic carbon concentration CB — ^- °°
and
CBO - CB =
where C...- = initial concentration of TOC
oU
C_, = concentration of TOC at time t
D
P, = partial pressure of ozone in the gas phase
H = Henry's law constant
a = ratio of stoichiometric constants c/b,
which come from
03 + hv ^-c R- C59]
kn(B)
B + bR'-2 ^- products (C02 + H20) [60]
-rB - -d CB/dt = koCRmCBn [61]
b) When C_, ^ 0
£>
ln(CBN/CB) = k2PA m/1t, for n = 1 [62]
CB1~n - CBN 1~n = (n-l)k2PAm/1t, for n ^ 1 [63]
where CB = CBN/ at t = tR
The order 1 in the last equation is derived from the
suicidal reaction for free radicals
-rR,S ° kS CR
For an intermediate range of CB concentration, the
following expression was applied,
-rB = - d CB/dt = k C/ = k'VP P^ C^ [65]
48
-------
where k1 = f(luv, T, pH, H)
Vg = gas superficial velocity
PA = Io9arithmic mean of ozone partial pressure in the
reactor and
luv = Total output of UV lamp (W) r n
Reactor wet volume (L) LbbJ
Based; on the experimental data, the authors obtained n=l,
q = 1.5 and p = 1.5 for the O3/UV of TOG in hospital wastewater;
Equation [65] thus illustrates the dependency of rate on UV
intensity, ozone input rate, and mass transfer efficiency.
In the next chapter expressions will be developed which
utilize and combine various features of the previous treatments,
in order to describe the kinetics of ozone/UV-induced reactions
and allow prediction of reaction rates from specified values of
the independent variables.
CONCLUSIONS FROM LITERATURE SURVEY
a) Ozonation has been used as a disinfectant of drinking
water for decades, without report of serious negative
consequences.
b) Ozone decomposes fairly rapidly in water, but there is
not general agreement as to the kinetics or mechanism.
Hydrogen peroxide and oxygen were found as final
products and hydroxyl radical (-OH) as an intermediata
The decomposition rate increases with increasing pH,
up to about pH 13.
c) The rate of oxidation of organic compounds increases
under conditions which favor decomposition of ozone,
d) The active species formed upon ozone autodecomposition
in water is the hydroxyl radical.
e) Reactivity of the hydroxyl radical is considerably
faster and more non-specific than ozone itself.
f) Ozonation products include carboxylic acids and
aldehydes predicted by cleavage of double bonds.
Phenols yield polyhydroxy compounds which then degrade
to smaller non-aromatic acids. Hydrogen peroxide has
also been found. 4,4'-dichlorobiphenyl gave 4-
chlorobenzoic acid.
49
-------
g) Organohalogen is converted to chloride but not
necessarily at the same rate as substrate disappears
providing there is a site for attack which does not
include the chlorine-bound carbon atom.
h) Ultraviolet light accelerates the decomposition of
ozone and therefore may provide a specie more
reactive than ozone.
i) Photolysis of dry ozone in'the vapor phase gives a
quantum yield of two; greater if water vapor is
present. Photolysis products are oxygen atoms and
hydroxyl radicals.
j) In aqueous solution hydrogen peroxide is a product
of ozone photolysis, probably due to combination of
oxygen atoms with H20. Half of the oxygen in H202
comes from ozone, half from water.
k) Hydrogen peroxide also photolyzes in water to produce
hydroxyl radical.
1) Treatment of phenol with ozone or H2O2/UV each gives
an initial increase in the chloroform formation
potential.
m) An initial increase in trihalomethane formation
potential, followed by a decrease is observed upon
ozonation of natural water.
n) Rates of aqueous photolytic ozonations are ten to one
thousand times faster than ozonation alone.
o) Mechanisms postulated for the ozone/UV reaction are:
1) Simultaneous interaction of substrate, ozone, and UV
2) Photolysis of substrate to free radical which
reacts with ozone.
3) Photolysis of ozone to form active species.
4) Photolysis of water to form the transients -H and
•OH
5) Excitation of substrate which then reacts with
ozone.
6) Excitation of ozone which then reacts with
substrate.
50
-------
p) Ozonolysis generally results in a plateau in TOG
representing refractory species while photolytic
ozonation will destroy these species.
q) Molecular oxygen is expected to play an important
role, not as the active species, but in secondary
reactions with intermediate products.
r) Water quality will play an important role in
determining the rate of an ozone or ozone/UV reaction,
and it would therefore be necessary to study the
reaction in natural water as well as in purified water.
s) Photonucleophilic displacement reactions result in the
conversion of aromatic organohalides to phenols and
halide ion.
t) Models for simultaneous ozone mass transfer and
chemical reaction indicate that ozone dose rate
(amount of ozone applied per unit time per unit
volume of liquid to be treated) would be the correct
parameter to use in the rate expressions for ozone-
related reactions. Depending on the efficiency of
the mass transfer process, there may be an efficiency
coefficient imbedded in the rate constants determined
using ozone dose rate as the intensity factor.
51
-------
SECTION 5
CHEMICAL KINETICS OF OZONE/ULTRAVIOLET-INDUCED
REACTIONS OF ORGANIC COMPOUNDS IN WATER
EXPERIMENTAL PROCEDURES
The Quartz Photochemical Reactor used in these studies was
cylindrical with a somewhat rounded bottom, and approximately 10
inches in height by 6 inches wide, into which 3L of sample was
placed (Figure 11). It has four baffles on the walls to aid
mixing, a six-bladed paddle impeller turned by a stirring motor
at 300-500 rpm, and several ports in the lid for introduction
of ozone, exiting of off-gas, taking of samples, etc. The gas
stream from the ozone generator was run through a modified
liquid chromatography UV absorbance detector where the concen-
tration of ozone in the inlet stream could be measured spectro-
photometrically. After being bubbled into the bottom of the
reactor and up through the liquid, ozone in the off-gas was
measured as above, then destroyed thermally. The entire quartz
reactor could be lowered into a chamber surrounded by UV lamps.
A description of the reactor itself and other experimental
details are presented more thoroughly in Appendix B.
During a typical kinetic run in purified water, the
desired ozone dose rate was equilibrated with the UV lamps on,
then a concentrated aqueous solution of the substrate was added
through a port in the reactor head. Samples were withdrawn as
a function of time and analyzed by gas chromatography. Ozone
in the liquid was measured iodimetrically, and the UV intensity
reaching the solution in the reactor was measured by ferriox-
alate actinometry.
Total Organic Carbon, where desirable, was measured by a
Dohrmann low-level DC-54 and trihalomethane potential was
measured by chlorinating the sample at a given level and then
measuring trihalomethanes after a specified waiting period
(3-8 days) at 26°C.
Because of its relative insolubility, hexachlorobiphenyl
(HCB) was dissolved in methanol, 18 yL of which was spiked into
the reactor to give the desired concentrations.
52
-------
VSPARGER-TOP
03/02 Inlet
r
'?"
I
T
-STIRRER
PTFE 0-RING
f
Excess 03/02 Outlet
-PYREX SAMPLING
TUBE
r\
GASKET
- QUARTZ BELL JAR
BAFFLE
IMPELLER
SPARGER
Figure 11. NTSU photochemical reactor,
53
-------
During the kinetic runs using lake or river water, the
sample was placed in the reactor with no ozone flow or UV, then
the ozone and UV turned on at the beginning of the experiment.
Sampling and analysis was as above.
RESULTS AND DISCUSSION
The purpose of the kinetic analysis is to provide a general
expression from which, for a given ozone dose rate and UV
intensity, the disappearance curve for substrate may be predicted.
This can in turn be used to optimize the ozone/UV process for a
given application.
The methods of kinetic analysis and the resulting rate
equations will be treated below, after which a more qualitative
and comprehensive discussion of the results will follow.
The kinetic results for the compounds studied can be
grouped into three classes:
1) Those which are first order in substrate, specifically
chloroform, bromodichloromethane , and tetrachloro-
ethylene.
2) Those of complex order, in this case hexachlorobiphenyL
3) Non-homogenous systems; trihalomethane precursors
(TTHMP) and total organic carbon (TOC) .
Since a different approach to the kinetic analysis was utilized
in each case, they will be discussed in separate sections below.
The generalized starting point, however, is the following.
The rate expression for disappearance of substrate is
assumed to be the sum of the rates of the individual reactions
taking place. For simplicity, the following reactions were
assumed:
1) Removal of the substrate from solution due to purging
by the sparge gas O^/C^; a first order process.
-If • S ' [67]
where S = substrate, and
k = first order rate constant
54
-------
2) Photolysis of the substrate in the presence of UV
radiation;
at = kp1 S [68]
where I = intensity of ultraviolet radiation, in W/L
3) Ozonolysis of substrate
_ dS - k T n 1 c<;d
dt V°3J£S [69]
where 0^ = ozone concentration in the liquid
4) Ozone/UV reaction
d£ Tenfo9
dt ~ u D S [70]
where D = ozone dose rate, in mg/L-min
The first three rate expressions should be correct
independently of the fourth, that is, for example, the
ozonolysis reaction should proceed at an instantaneous rate
which is proportional to some power of the ozone concentration,
and should therefore be measureable in a separate experiment.
This assumption is valid when the "ozone-only" reaction is not
so fast that the solution becomes depleted of ozone and the
rate of reaction is so mass transfer limited that reaction takes
place at the bubble surface. The same assumption of indepen-
dence is used for the photolysis and purging reactions. On the
other hand, there may be several reactions lumped into the
fourth term, particularly where the substrate is a good UV
absorber. Reaction of substrate with the active specie
generated from the photolysis of ozone is probably the major
reaction which takes place, although, as discussed later,
reaction of photoexcited substrate with ozone is a possibility.
The reaction of substrate with photoexcited ozone would also
appear in the fourth term, should it occur. The general
approach used was to evaluate the first three terms from
independent experiments, then to subtract their contribution
from the experimentally determined total rate of the ozone/UV
experiment. The "remainder" of the rate was then empirically
fit to the form kuIeDfS?. As mentioned in the preceding section,
the use of the "ozone dose rate" in the kinetic expression was
55
-------
previously suggested by Eisenhauer (142) and indirectly by Lee
and co-workers (110) . The implication that it should be correct
for ozone without UV in the case of dilute solutions of sub-
strate in pure water in a serai-batch reactor is not so obvious,
since the equilibrium value of dissolved ozone concentration is
proportional to the concentration of ozone in the gas stream,
not the absolute amount. However, the greater quantity of a
gas of a given concentration that is present, the less that gas
stream is depleted. In the case of the ozone/UV experiments in
this investigation, very little (2-8% of the incoming ozone
concentration) ozone was found in the off -gas from the reactor,
and ozone in the liquid phase was below the detectable limit in
almost all cases, so that dose rate was the logical parameter.
Details are explained by compound in "the following section.
Because of the complexities of the processes and the
inhomogeneity of the substrate, the above procedure was not used
to analyze the trihalomethane precursor destruction data.
Analysis of that data will be discussed in a separate section
below.
Kinetic Analysis of Ozone/UV Reactions Which are First-Order
in Substrate
In the case of chloroform (CF) , bromodichloromethane (BDM) ,
and tetrachloroethylene (TCE) , the sum of the terms representing
the individual reactions (eq. 67-70) can be condensed to a
pseudo-first-order rate constant since the overall reaction
rate during an ozone/UV experiment is found to be first-order
with respect to (directly proportional to) substrate, as were
the individual processes [67] through [69] ,
- If =
= EK S = KS .
i *
This pseudo-first-order rate constant, K, is itself a sum of
pseudo-first-order rate constants, the relative values of which
represent the relative importance of the processes comprising
the total reaction.
I. Purified Water as the Reaction Medium
Details of the kinetic analysis, being necessarily
mathematical, are given in Appendix C, Section I. The results
are discussed in Section III below.
56
-------
II. Lake Water as the Reaction Medium
In an actual drinking water treatment situation the water
will not be highly purified. The presence of substances other
than the substrate of interest will decrease the rate at which
that substrate is destroyed, due to competition for the active
species. Hoigne1 and Bader have published several papers (40,44,
45,46) in which they discuss the effect of TOC, ammonia, bicar-
bonate, and carbonate on the destruction rate of micropollutants
by ozone, particularly in the high pH range, where they give
evidence that hydroxyl radical is the active specie generated by
the decomposition of ozone. Since a similar effect is expected
with ozone/UV, the reaction rates of the model compounds were
also studied in a natural water matrix, Lewisville Lake Water
(LLW) .
This water was partially characterized as follows:
pH 8.1
Suspended Solids <10 mg/L
Total Dissolved Solids 130 mg/L
Chemical Oxygen Demand 10 mg/L
Total Organic Carbon 4 . 3 mg/L
NH3/NH* (as N) 120 yg/L
CO=/HCOJ 1 mM
If the expression of Hoigne' and Bader (45) is used to
calculate the decrease in reaction rate in going from pure water
to Lewisville Lake Water, the rate changes by the ratio of the
"competition values," ft. Since
«« ,. • = (nVkj'1 Z k. S. , [72]
M, matrix a M i i i
where M is the test compound, n's are fractional efficiencies
characteristic of pH and compound M, k's are PFO rate constants
and S. is the concentration of component i, the ratio is
given By
kM SM [73]
M. Pure water =
a S k S
"Mf LLW i i i
Evaluation of this expression using reference 45 and the 'above
data on Lake Lewisville water gives
57
-------
kCHC!3 = 1.9 X 107 M~1sec l , [CHC13] = 8.5 X 10"7 M
kTOC = 1 X 109 M'1 sec'1 , [TOC] = 4.3 X 10~5M
kCC>3 = 3 X 108 M"1 sec'1 , [C0~ ] = 7 X 10~6 M
ku_ - = 1.7 X 107 M"1 sec'1 , [HCO~] = 10~3 M
ilv_Uo j
kNH/NH+=1.3 X 10 7M"1 sec'1 , [NH/NH]= 6.5 X 10~6
3. , 3 . M
and
ft
'M,pw _
[,LLW ~ CHC13 + TOC
16.15 J..6XL01
1.6X101 + 4.3X104 + 2.1X103 +1.7X104 + 8.5X101 6.2XL04
= 2.6 X 10~4.
Since the water used in the purified water experients is not
strictly pure but contains typically 0.2-0.4 mg/L of TOC, this
factor of lO'* difference in rates is not to be expected experi-
mentally, but rather
, purified water 1.6X101 + 3X10~6(109) 3 X 103 1
I LLW 6.2 X 104 " 6.2X104 " 20
where an average figure of 0.3 mg/L has been used for the TOC of
purified water. This calculation shows that if the active specie
were hydroxyl radical, the rate could be expected to slow by a
factor of twenty in going from purified water to this natural
water matrix. Although the identity of the ozone/UV active
specie is not known, all evidence in this study indicated a very
reactive free radical species, so that the possibility of a
similar matrix effect existed. In view of this large difference
in reaction rate, it was necessary to perform kinetic studies in
natural water and to use that kinetic analysis as the basis for
the engineering calculations. Appendix C, Section II describes
the kinetic analysis on a compound-by-compound basis, and the
results are discussed in Section III, below.
58
-------
III. Discussion of Kinetic Results for Tetrachloroethylene,
Chloroform, and Bromodichloromethane
Because of the complexity of the kinetic analyses described
in Appendix C it is difficult to visualize the differences and
similarities between rate expressions and to make comparisons
between compounds and matrices. The choice of model compounds,
based at the time of choosing on the interest in those compounds
as potential pollutants, was a fortunate choice. Hexachloro-
biphenyl, described in another section of this report, is an
aromatic compound. Tetrachloroethylene (TCE) contains a double
bond but is not aromatic, while chloroform (CF) and bromo-
dichloromethane (BDM) contains no double bonds. The latter two
compounds do, however, have one important dissimilarity: BDM
absorbs ultraviolet radiation in the region of the important
mercury lines while CF does not.
The ozone/UV reactions of TCE, CF, and BDM are all first
order in substrate, which makes them somewhat easier to compare,
since their rate equations all have the same general form,
RATE = -S = -~=KS .
at
which states that the rate of reaction is proportional to the
concentration of substrate present. For a given substrate
concentration, then, the relative rate of reaction of two com-
pounds is the same as the relative size of the rate constants.
Since the functional dependence of the rate constants on ozone
dose rate and UV intensity has been determined in the previous
section, it is instructive to compare the form of those rate
constants as well as the contribution of the various reactions.
Table 3 summarizes the general form of the rate equation and
the more specific variations which each compound takes in the
two matrices, as determined in Appendix C.
Points which should be especially noted are:
1) Increasing the UV intensity has less effect on TCE in
lake water than in purified water.
2) The ozone/UV term of chloroform in purified water has
a very weak UV dependency. Upon changing to lake
water that UV dependence drops to zero.
3) The odd orders of dose rate dependence for chloroform
indicate that its reaction is kinetically very
complicated. This point is discussed further in the
products section of this report, where further
evidence for a chain reaction is given.
59
-------
TABLE 3. KINETIC FORM OF PSEUDO-FIRST ORDER RATE CONSTANTS FOR
TETRACHLOROETHYLENE, CHLOROFORM, AND BROMODICHLOROMETHANE
(a)
Compound
GENERAL
TCE
CF
BDM
(b)
Matrix
PW
LLW
PW
LLW
PW
LLW
Purging
j^
purge
j^
purge
j^
purge
purge
k
purge
k (g)
purge^5
k
purge
(c)*
Ozonolysis
k D ™ -
o eff
k D ^i
o eff
koDeff
koDeff
k D ff
o eff
k D ,f
o eff
k D ff
o eff
(d)
Photolysis Ozone/UV
k IX k I*DZ
P u
k I k ID
p u
k i°-670 k i°-419
V u1 D
k i k 2-266 !
kpZ u1 D
k I k D°-562
P u
-0.466
k I k I D
P u
knU) kj0.481
p u
Notes
•5(f)
(h)
(j)
a) TCE=tetrachloroethylene, CF=chloroform, BDM=bromodichloro-
methane
b) PW=purified water, LLW=Lewisville Lake Water
c) In which no UV is involved
d) In which no ozone is involved
e) Converted to this form from that given in the text using the
experimental relationship [03]. = 2.46 X 10~4D for ozonolysis
experiments. v
f) In the curve fit only the ozon«/UV term was used because of
its greater magnitude compared to the other terms
g) The small correction term, kc, is not included in this table;
see text.
h) Provided D is below the cut-off point; see Appendix C
i) Only two data points; assumed first order, checked by
prediction of k . See Appendix C.
j) Restriction of correspondence of terms to different
chemical reactions relaxed. See Appendix C.
60
-------
TABLE 4. REACTION RATE CONSTANTS k^ FOR INDIVIDUAL TERMS IN RATE
EXPRESSION OF TETRACHLOROETHYLENE, CHLOROFORM, AND BROMODICHLORO -
METHANE (a)
Compound Matrix Purging
]r
purge
TCE PW
LLW
CF PW
LLW
BDM PW
LLW
LLW
1.
9.
1.
9.
1.
2.
-7.
24X102
96X103
5 X103
54X104
43X103
65X104
39X103
Ozonolysis
ko
8
3
8
2
3
4
.66X10~2
.68X10~2
.0 X10~3
.56X10~3
.45X10~3
.86X10"3
Photolysis Ozone/UV
k k
P u
5
7
2
4
5
2
8
.68X101
.86X102
.5 X102
.58X103
.62X102
.99X102
.95X102
6
1
4
9
6
5
.54
.78X101
.78X101
.18X103
.56X102(C)
.04X103
Notes
(b)
(d)
(e)
a) Units of time - minutes, of ozone dose rate - mg/L min, of W
intensity - W/L. Powers of units consistent with Table 3.
b) Rate constants for purging, ozonolysis, and photolysis are
evaluated from separate experiment; the ozone/UV rate
constant given is for the total reaction.
c) Does not include k = 8.78 X 10~~ . See Appendix C.
C
d) Determined from separate experiments.
e) Empirical expression to be used for calculating 03/UV runs
only (see Appendix C).
61
-------
TABLE 5. CONTRIBUTIONS TO PSEUDO-FIRST-ORDER RATE CONSTANTS
BY THE VARIOUS COMPONENT TERMS.
(b) (c)
Compound Matrix Purging Ozonolysis
TCE PW 1.24X102 0
LLW 9.96X103 0
CF PW 1.5X10~3 0
LLW 9.54X104 3.1X10~5
BDM PW 1.43X103 0
LLW 2.65X104
(d)
Photolysis
1.
2.
5.
9.
1.
5.
14X10"1
67X10~2
oxio"3
16X104
12X102
98X103
(d)
Ozone/UV
7.
2.
7.
1.
1.
1.01
03X10"2
isxicT1
95X10"3
osxio'1
28X10~2
(f)
1
2
9
1
1
Total
1.13
.07X10"1
.20X101
.85X103
.20X101
.90X102
(e)
a) All values in minutes . Values of reaction conditions
assumed are I = 0.20 W/L, D = 0-775 mg/L min.
b) TCE = tetrachloroethylene, CF = chloroform, BDM = bromo-
dichloromethane
c) PW = purified water, LLW = Lewisville Lake Water
d) From line 6 of Table 4
e) From line 7 of Table 4
f) by difference
62
-------
TABLE 6. PERCENTAGE CONTRIBUTION TO PSEUDO-FIRST-ORDER
RATE CONSTANTS BY THE VARIOUS COMPONENT TERMS(a)
Compound Matrix
TCE PW
LLW
CF PW
LLW
BDM PW
LLW
Purging
1.9
9.3
0.7
9.7
1.2
1.4
Ozonolysis
0
0
0
0.3
0
0
Photolysis
10
25
2.3
9.3
9.3
31
o3/uv
89
66
97
81
90
67
a) calculated from the values in Table 5
TABLE 7. RATIO OF RATE CONSTANTS IN PURIFIED WATER TO THOSE
IN LAKE LEWISVILLE WATER
Compounds Purging Photolysis °3/uv Total
TCE 1.2 4.3 14.4 10.6
CF 1.6 5.5 26.8 22.3
BDM 5.4 1.9 8.4 6.3
63
-------
4) The negative order of the BDM-Oo/UV reaction with
respect to UV intensity is unlike any other compound
studied, and may indicate competing reactions.
5) Not shown in Table 3, but indicated earlier is the fact
that for bromodichloromethane in purified water,
increased ozone dose rate past a certain value results
in no further increase in rate constant. This cut-off
value of D seems to double when the UV intensity is
doubled, indicating a direct relationship.
The rate constants, K.J_ , are collected in Table 4 in the
units appropriate to Table 3. By tabulating values of the
pseudo-first-order rate constants representing the individual
reactions, evaluated for a given set of reaction conditions, the
relative contributions from each term to the total can be seen.
This is done in Table 5 for values of ozone dose rate = 0.775
mg/L min and UV intensity of 0.2 W/L. The chosen values of I
and D represent the middle of the experimentally investigated
range of these parameters. In the last column in Table 5 are
shown the total PFO rate constants for comparison with the
contributed parts. The relative contributions of the various
terms is made more clear in Table 6 where the percent of the
PFO rate constant contributed by each reaction is shown.
Finally, the results are shown in another way in Table 7, where
the ratios of the PFO rate constants for the various reactions
in purified water compared to Lake Lewisville water are tabu-
lated. These are the factors by which the reaction rate
decreases when going from purified water to Lewisville Lake
Water.
Several points concerning the results tabulated in Tables
4-7 are worthy of note:
a) Because of the destruction of ozone by UV there was no
ozone present in solution in 90% of the purified water
experiments. Only for the highest dose rates with the
lowest UV intensities was ozone found in solution, and
then only in amounts equivalent to an effective dose
rate of 0.05 mg/L min and less. In the lake water
experiments the natural components of the water absorbed
some of the UV so that the actual intensity available to
an effective dose rate of only 0.13 mg/L min, approxi-
mately equal to the lowest actual dose rate used in any
of the experiments. Because of this, the ozonolysis
contribution to these reactions is essentially zero,
as seen in Table 6.
64
-------
b) Although TCE and BDM are quite different chemically and
have very different rates of reaction, they have
approximately the same ratio of purging and photolysis
rate to the total PFO rate constant, as shown in
Table 6. Chloroform, which does not appreciably absorb
in the UV region naturally has much lower photolysis
contribution, ^2% in PW compared to 9-10% for TCE and
BDM. However, because of the rate of the ozone/UV
reaction, CF is intermediate in reaction rate between
TCE and BDM.
c) Rather than experiencing a "salting out" effect, all
three compounds appear to purge more slowly from lake
water than from purified water (Table 7). This may
indicate a tendency of the compounds to complex in
some way with the constituents of the lake water.
d) Next to purging, photolysis of the model compounds
appears to slow down the least in going to a natural
matrix, as seen from the ratios in Table 7. This is as
expected, since the UV absorbance of LLW is not large.
It should be noted that ratios for ozonolysis, k0; can
be calculated from the data in Table 4 and give lower
values than those listed for photolysis in Table 7.
The k values listed in Table 4, however, are based on
very limited data so that the ratio must be regarded
carefully.
e) The CU/UV reaction rate of chloroform is seen in Table
7 to decrease by the greatest factor upon the addition
of the natural matrix. If indeed a chain reaction is
involved in the destruction of chloroform, introduction
of the natural matrix could not only compete for the
original active species but also for the chain carrier.
f) In purified water the total reaction of TCE is about
five times faster than chloroform which is in turn
twice as fast as bromodichloromethane. In natural
water, TCE is about five times as fast as BDM which is
twice as fast as chloroform (Table 5).
g) The decrease in reaction rate of the model compounds in
going from purified water to natural water is by a
factor of 6 - 22. It should be noted again here that
the rate in purified water is not the same as the rate
in pure water, as calculated in section II above.
h) If the substrate is a UV-absorber, the photolysis con-
tribution to the total rate of an ozone/UV reaction is
likely to be appreciable, while, as mentioned above,
the ozonolysis contribution may be quite negligible.
65
-------
i)
Most importantly, for a given does rate, ozone/UV is
considerably faster than ozone alone, as shown in the
following table of relative rates, calculated for the
same values of I = 0.2 W/L and D = 0.775 mg/L min as
in the previous example.
OZONOLYSIS ONLY
PHOTOLYSIS ONLY
OZONE+UV
TCE/PW
/LLW
CF/PW
/LLW
BDM/PW
/LLW
(D = 0.775)
.059
.27
.028
.20
.022
.20
(I = 0.20)
.10
.25
.023
.093
.093
.31
1
1
1
1
1
1
It is seen that even in the least favorable case, the
combination of ozone and UV is four times as fast as
either ozone or UV alone, and in the most favorable
cases is fifty times as fast as ozone and forty times
as fast as UV alone. It is also seen that there is a
striking similarity between all three compounds in the
decreasing efficiency of 03/UV relative to ozone when
going from purified water to a natural matrix.
j) Finally, although previously unanticipated, UV alone
suggests itself as a treatment method which has the
specificity for certain undesirable micropollutants
which ozone/UV seems to lack. Since the major expenses
of an ozone/UV system are the ozone generator and
power to run it, plus the contactor, ultraviolet
radiation could prove to be very cost-effective for
the compounds for which it was specific. This point
will be discussed in more detail later in this report.
The effect of the ozone dose rate, and UV intensity on
the pseudo-first-order rate constant, and thus on the
reaction rate is shown in Figures 12-15 for the entire
range of experimental values. The situation for
tetrachloroethylene (Figure 12) is seen to be first-
order in dose rate. In Figure 13 the lake water por-
tion of the plot is magnified to show its behavior as
well as for comparison with the other two compounds.
In Figure 14, the results for chloroform, are seen the
concave curves which indicate the reaction to be
66
-------
4.0 r-
0.375 W/L
PURIFIED WATER
LEWISVILLE LAKE
0.189 W/L
0.091 W/L
WATER
0.375 W/L
0.5
OZONE
DOSE
1.0 1.5
RATE, mg/L - min
Figure 12.
Pseudo-first-order rate constants for ozone/UV
destruction of tetrachloroethylene.
67
-------
C/5
z
o
o
0.20 r
0.15
0.10 -
0.0
0.189W/L
PURIFIED WATER
0.375 W/L
0.091 W/L
LEWISVILLE LAKE WATER
0.5 1.0
OZONE DOSE RATE, mg\L-min
1.5
Figure 13. Pseudo-first-order rate constants for ozone/UV
destruction of tetrachloroethylene, detail.
68
-------
0.20 r-
.189 W/L
.084 W/L
PURIFIED WATER
LEWISVILLE LAKE WATER
.375 W/L
189 W/L
0.5 1.0
OZONE DOSE RATE, mg./L-min
Figure 14.
Pseudo-first-order rate constants for ozone/UV
destruction of chloroform.
69
-------
0.20.-
H 189 W/L
PURIFIED WATER
.375 W/L
. 084 W/L
LEWISVILLE LAKE WATER
.375 W/L
.189 W/L
0.5
OZONE DOSE RATE,
1-0
mg/L - min
1.5
Figure 15.
Pseudo-first-order rate constants for ozone/UV
destruction of bromodichloromethane.
70
-------
greater than first- order in dose rate in purified
water. Also obvious is the much greater difference
between the rate in pure water and lake water,
compared to TCE.
Shown in Figure 15 is the most complicated case
discussed so far, that of bromodichloromethane . The
cut-off dose rate at approximately 0.56 and 1.07 mg/L
min for one and two lamps, respectively, is most
obvious in this plot, while the negative dependency of
the lake water rates on the UV intensity can barely
be discerned.
IV. The Effect of pH on the Reaction Rates of Tetrachloro-
ethylene, Chloroform, and Bromodichloromethane.
Because of the great amount of data collection and analysis
necessary for the completion of the previous and subsequent
sections, the effect of pH on the rate of ozone/UV destruction
of tetrachloroethylene, chloroform, and bromodichloromethane
was not investigated comprehensively enough to allow a kinetic
analysis. There is, however, interesting and important infor-
mation which can be gathered from the pH studies which were
performed.
A. Tetrachloroethylene — Destruction of tetrachloroethylene
by ozone is fairly rapid compared to chloroform. Comparison of
T(3) (95% removal) for pH 4, 6.9, and 10 is shown in Figure 16,
where it is seen for example that for an ozone dose rate of 0.8
mg/L min, the reaction at pH 10 is five times as fast as that at
pH 4. This order is consistent with the suggested (42) hydroxyl
radical mechanism at high pH. On the other hand, the ozone UV
process takes four times as long at pH 10 as at pH 7 which is
1.5 times slower than pH 4. The data is shown in Table 8.
TABLE 8. ELIMINATION TIME, T , FOR TETRACHLOROETHYLENE BY
OZONE AND OZONE/UV FOR VARIOUS pH VALUES
Process pH 4 pH 7 pH 10
Ozone
Ozone/UV
5000 sec
53 sec
2750 sec
89 sec
450 sec
395 sec
(a) corresponding to 95% removal
(b) ozone dose rate =0.8 mg/L min, UV intensity = 0.375 W/L
71
-------
10,000
8,000
6,000
T (3),
sec.
4,000
2,000
pHIO
I
I
I
I
02
0.4
J
1.4
Figure 16.
0.6 0.8 1.0 1.2
OZONE DOSE RATE, mg/l-min.
Tetrachloroethylene - effect of pH on treatment time
(95% removal from 100 yg/L ozone only (no UV)).
72
-------
B. Chloroform—Chloroform behaves in much the same manner
as tetrachloroethylene in the response of its reaction rate to
pH. Figure 17 shows the effect of pH on the ozonolysis of
chloroform. It is seen that at the end of 100 minutes, about 18%
of the chloroform has been removed by ozone at pH 7, while a
similar dose rate at pH 10 has removed 56%. On the other hand,
similar doses of ozone with UV leave 12% remaining after 120
minutes at pH 9.8 while at pH 6.5, less than 3% remains after 15
minutes (Figure 18 )•
C. Bromodichloromethane—Figure 19 shows the removal due to
ozonolysis of bromodichloromethane from solution at pH 3.2, 6.5
and 10. It is seen that BDM, like TCE and CF, is removed much
more quickly at pH 10. Experiments in which no UV or ozone was
present showed that there was some basic hydrolysis of BDM
taking place at pH 10, but it is not fast enough to contribute
appreciably to the rate. Also as in the case of TCE and CF, the
removal of BDM by ozone/UV is enhanced by low pH, as shown in
Figure 20.
Thus all three of the model compounds discussed thus far
react similarly in that high pH speeds the ozonolysis process
while the ozone/UV process is fastest at low pH. This similarity
is interesting in view of the markedly different rates of
reaction and the unusual differences in the rate equations of the
three compounds. This pH behavior may contain important infor-
mation about the identity of the primary active species, since
the behavior is constant for different substrates. This point
will be discussed further after additional data is introduced.
Kinetic Analysis of the Ozone/UV Destruction of Hexachlorobiptenyl
The study of the model compound 2 ,2' ,4,4',6,6'-hexachloro-
biphenyl presented unique difficulties. First of all, its very
low solubility made it necessary to perform the experiments using
an initial concentration of 60 yg/L. This fact, coupled with
the compound's electron capture response made it necessary to
perform a concentration step on each of the individual samples,
thereby affecting the precision of the measurements. Studies
were performed using initial concentrations lower than 60 ug/L to
ensure that the solubility limits had not been exceeded and that
the same rate law was applicable.
The fact that HCB was very slowly soluble in water made it
necessary to dissolve the compound in methanol and then to spike
water with the concentrate to obtain the desired initial concen-
tration. This procedure introduced 6 yl of methanol per liter
of solution which could alter the reaction rate considerably in
purified water. That quantity of methanol is equivalent to
4.7 mg MeOH per liter, as compared with a TOC value of 4.3 mg/L
for the lake water.
73
-------
PH7 3.12
D pHIO 2.53my03
'linin
0
In
[CHCI3]
[CHCI3]
-2
10 20 30 40 50 60 70 80
TIME{MIN.)
I
90 100 110 120
Figure 17. Effect of pH on ozonolysis of chloroform.
-------
pH6.5 0.8mg/l-min
pH9.8 0.7mg/l -min. 03
j
10 20 30 40 50 60 70 80 90 100 110 120
TIME (MIN.)
Figure 18. Effect of pH on photolytic ozonolysis of chloroform-
75
-------
0.6
[CHCI2Br]
[CHCI2Br]
0.4
0.2
0
= 3.2
pH = 6.5
20
OZONE DOSE RATE = 1.4mg/l-min.
pH = IO
1000 2000 3000 4000
5000 SEC.
40
TIME
60
80 MIN.
Figure 19. Effect of pH on BDM destruction by ozone.
76
-------
O 0.07mg 03/l-min, pH= 4
Q 0.10mg 03/l-min,pH = 8
A QOTmg 03/l-min, pH=IO
1000 2000 3000 4000 5000 6000
TIME, SEC
Figure 20. Effect of PH on ozone/UV destruction of BDM.
77
-------
The HCB reaction was of apparent kinetic order greater than
one for both photolysis and ozone/UV. High apparent orders are
amplified by high initial rates which slow tremendously as the
reaction proceeds. Thus, only with the highest dose rates and
UV intensities used did destruction ever reach 90% in the
eighty-minute reaction time which was allowed. It was because of
this slowness in the second half of the reaction that HCB has not
been included in the process design calculations, since it would
obscure information about the other substrates, all of which
were closer together in their reaction rates. The kinetic
results are of interest, however, even if not completely under-
stood, and are included for completeness. This portion of the
work is presented in more detail by Huang (161) , whose results
are summarized in Appendix D and briefly below.
I. Lake Water as a Reaction Medium
The direct ozonation of HCB in lake water is slow, just 50%
of the 60 yg/L initial concentration being destroyed in one hour
at the highest ozone dose rate used, 4 mg/L min. The disappear-
ance curves (Figure 21) have an unusual appearance, the rate
being fairly rapid at first, then slowing down considerably
within the first five or ten minutes. In all figures pertaining
to HCB, the points are actual data points while the solid lines
are curve fits from the empirical equations derived. Runs were
conducted with no UV or ozone and with sparging of the reactor
with nitrogen to determine if the drastic change in rate was due
to some phenomenon such as adsorption, but no disappearance was
seen. The initial rate was found to be directly proportional
to dose rate.
The photolytic ozonation of HCB was considerably faster than
ozonation, but still slow, taking 1 hour to reach 90% destruction
at the highest dose rate (4 mg/L min) and a UV intensity of
0.25 W/L (Figure 22). After the photolysis rate (proportional
to the 1/2 power of the UV intensity) was subtracted out, the
ozone/UV portion of the rate expression was found to be directly
proportional to ozone dose rate and UV intensity, and propor-
tional to the square of the concentration of substrate present.
It is worth noting that at long reaction times, the photolytic
ozonation rate is in many cases very similar to the photolysis
rate.
II. Purified Water as the Reaction Medium
«
The same form of the photolysis equation was used for HCB
in purified water as for lake water, but the rate constant (and
consequently the rate) was about 1/3 larger (faster). Sub-
traction of the photolysis contribution from the total rate of
disappearance of HCB gave an expression which was 1^ order in
dose rate, first order in UV intensity, and proportional to the
cube of the substrate concentration. It should be stressed again
78
-------
I.U
[HCB]
[HCB]
u o
.5
0
(
jg-l | — -, , 1 , , , ,
fro
A
— ^
D 0
n
- • ° 8
•8 A 0 0 0
A 00©
- n O
» D A A O
- D A D A A 0
- •••; «.n - .
^ ^ ^ n i— i
... ^ 0 0
•
~~ D (mg/min-L)
0 1.16
A 2.19
D 3.00
• 4.04
, I , 1 , 1 i 1 . 1 i 1 > 1 i 1
D 5 10 20 30 40 50 60 70 80
TIME (min)
Figure 21.
Effects of dose rate on the direct ozonation of
HCB in lake water.
79
-------
1.14
2.16
3.04
3.95
w/L
PREDICTED
PHOTOLYSIS
111
0 5 10
I
20
I ' I ' I
30 40 50
time , (min)
60
I
70
I
80
Figure 22.
Effects of dose rate on the photolytic ozonation
of HCB in lake water.
80
-------
that these are empirical fits to the data, useful for prediction,
and no mechanistic explanation can be offered, being totally
beyond the scope of the project.
HCB differed from the compounds previously discussed in that
elevation of the pH increased the destruction rate of both
ozonation and photolytic ozonation. It also differed in that
its destruction rates in lake water were very similar to those in
pure water and the two were, in some cases, indistinguishable in
the final stages. This may be due to the methanol which was
present in the pure water experiments, as discussed earlier.
These results are discussed in more technical detail in Appendix
D.
Kinetic Analysis of Ozone/UV Reactions with Nonhomogeneous
Systems; Trihalomethane Formation Potential and TOC;
The term nonhomogeneous systems in the context of this
report means that the substrate is not a chemically well-defined
species. More specifically", the systems of interest in this
study are trihalomethane formation potential (THMFP) and, to
the extent that it affects THMFP, total organic carbon (TOC).
So far, the entities responsible for THMFP have eluded
investigators, although several functional groups have been
discussed as contributors to the formation potential of a water.
The consensus seems to be that THMFP should be thought of as
sites, probably located on macromolecules such as those which
comprise humic and fulvic acids. This view is consistent with
the finding of certain functionalities which yield trihalo-
methanes upon chlorination. If this is the case, an important
issue is whether one can destroy the site without having to
remove or destroy the entire organic matrix. For this reason
it is important to monitor the TOC destruction in a process
which is designed to chemically destroy THMFP, even though
removal of the TOC may not in itself be a primary goal.
In preliminary experiments run in the reactor at Houston
Research, Inc., in Houston, Texas, and analyzed at North Texas
State University, it was indicated that photolytic ozonation
might actually produce THMFP as well as destroy it, an effect
which was also reported by Rook (15), Lawrence (32), by Riley
(162) for ozonation. These data are presented in Figure 23,
where it is also seen that the rate of destruction seems to be
roughly the same for ozone and ozone/UV down to approximately
S/S =0.5, but the asymptotic limit reached by ozone was quite
different from that reached by ozone/UV. Although this was in
agreement with similar previous claims for the ozone versus_
ozone/UV destruction of model substrates (103), it was not in
agreement with results obtained previously in this project for
81
-------
^
§
_J
H
Z
LJ
K-
0
Q.
FORMATION
UJ
z
J5
UJ
2
3
I
cc
1-
Q
,!_
2.10
2.00
1.90
1.80
1.70
1.60
1.50
1.40
1.30
1.20
1.10
1.00
.90
.80
.70
.60
.50
.40
.30
.20
.10
0
i
V
1
i
; '.\ _
't«\ ^*- "^-— ^
j^^^ ^ -^
f ^)\
V \ 02 only
\ \
\ • ._ O flmn/lmin^
1 * —---»- ^*a » liny* ^ IMIII /
^ 1 \
' -i ; o,+ uv (200 w)
\'9\
1« * " ........ . ^\ J. 1 l\ J f^R^\i m\
•\ • i 0,+ UV (65Ow)
', . '• \ o
',\ -. \ 03+ UV(650w),50c
'I'
;\\\
I \'A
*•--**.
A '"%-
V v^»
\ """*
^ ^
• \ V ^
\\
V ***••
\.
\,
\>......
*-'— ^4^~:~'. -J
1
0
15
30
60
TIME, (MIN)
90
120
Figure 23. Ozone and ozone/UV destruction of TKM precursors
Ohio River water (2/77)-
82
-------
Ohio River water as well as for 1 mg/L solutions of commercial
humid acid solution (Aldrich) , as shown in Figures 24 and 25
respectively.
Thus, it was decided to repeat these studies at lower ozone
dose rates which are more realistic for purposes of future
applications, and which also would be more likely to show
differences between the effectiveness of 0-> and 03/UV. For this
purpose, water from Caddo Lake in East Texas was chosen as the
test medium. This water represents surface water sources in the
south and southeastern parts of the USA where total organic
carbon levels reach high values and where THM values are expected
to be correspondingly high after chlorination. Caddo Lake water
at the time of our sampling showed a TOG value of 25 mg/L and a
DOC of about 12 mg/L.
Figure 26 shows the destruction of THM precursors in Caddo
Lake water with ozone doses of 0.14, 0.28 and 0.42 mg 03/L-min,
which correspond to 2.9, 5.8 and 8.8 ymoles/L-min. The essential
features of this data are as follows:
a) Under the conditions utilized, ozone initially destroys
high THM precursor levels in Caddo Lake water at a very
rapid rate. Initial slopes of the THM precursor
destruction curves for the 2.9 and 5.8 um 03/L-min runs
are 0.011 and 0.016 ymole THMFP/L-min.
b) Complete destruction of THM precursor levels is not
achieved in any of the runs. Indeed there appears to
be evidence for the formation of precursors as the
ozonation of the natural organics in the water proceeds.
Figure 27 is a plot of the data shown in Figure 26 after
subtraction of an exponential THMFP-destruction curve
obtained from the initial slopes of Figure 26. The
residual curves suggested a mechanism such as the
following, in which precursor P2 is formed by ozonolysis
of "natural carbon" C. P^ is precursor initially
present in the water, and presumably differs from P2
in chemical type and relative reactivity with ozone.
kl
p + o_ — = - ^- X X,Y = non-precursors
P2
Shown in Figure 28 is a comparison of one run in which a
dose rate of 0.42 mg 03/L min was used, with an 03/UV run at
83
-------
1.00
15
60
TIME (WIN)
90
120
Figure 24.
Ozone and ozone/UV destruction of THM precursors
Ohio River water (2/76).
84
-------
120
TIME IMIN)
Figure 25.
Ozone and ozone/UV destruction of THM precursors
humic acid (Ippm).
85
-------
O 0.14 mg 03/ I - min
A 0.28mgO I - min
D 0.42 mgO,/1 - min
0 15 30
60
90
TIME (MIN.)
150
180
Figure 26.
Normalized THM formation potential: Ozone destruc-
tion of THM precursors - Caddo Lake water .
86
-------
10L-
0.28mg 03/l - min
0.42mgO_/|-min
Exponential decay of initial
THM precursor (P )
Residual curve for formation
and decay of by-product
precursor (P_)
Figure 27.
90
TIME (MIN.)
Analysis of THMFP destruction curves - Caddo Lake
water.
87
-------
1.0
O 0.14mgO,/I-min; UV = 03I
A 0.42mgO-./l-min; no UV
15 30
60
90
TIME (MIN.)
120
150
Figure 28.
Normalized THM formation potential: Ozone destruc-
tion of THM precursors - Caddo Lake water .
88
-------
0.14 mg 03/L min and 0.31 W/L of UV radiation from medium
pressure mercury lamps. A comparison of the curves shows two
distinguishing features: (a) the 03/UV treatment destroys THM
precursor levels at a more rapid rate, and (b) no residual THMFP
remains in the case of the ozone/UV runs, whereas ozone alone
does not completely destroy THMFP after a net consumption of 930
mg of 03 (dose rate minus off-gas minus residual) during 180
minutes. It should also be noted that in this example, ozone
has produced additional precursor, as shown by the hump in the
destruction curve, whereas 03/UV did not, exactly the opposite
as was found for Ohio River water.
Following a decision to construct a reactor at NTSU, more
THMFP destruction runs were made using a single large water
sample. A more complete matrix of run conditions was used and
the data was accumulated as quickly as possible in order to
avoid deterioration of the water sample between runs. The
results of this matrix of experiments are shown in Figures 29 -
31. Experimental data is denoted by points while the curve fit
of the kinetic analysis is shown by the solid lines. Several
qualitative features of the data points stand out:
a) For a given UV intensity, destruction rate has an
obvious dependence on ozone dose rate.
b) There is clearly some sort of precursor generation at
the beginning of many of the runs. The size of the
error bars and the general smoothness of the data
attest to the reliability of the results, even with
the inherent imprecision of the THMFP analysis.
Secondary humps can clearly be seen in the following
runs:
I = 0.196 D = 6.36 t = 30
I = 0.096 D = 4.00 t = 45
Secondary humps have been ignored in the kinetic
analysis.
c) A general increase in THMFP destruction rate is seen as
UV intensity increases. There are, however, exceptions
to that trend:
1) I = 0.196, D = 6.36 > I = 0.40, D = 6.18
2) I = 0.196, D = 2.09 < I = 0.096, D = 2.23
In both of these cases, it can be seen that the run
which is slower than expected had to overcome a
greater initial buildup of precursor.
89
-------
FTTHMFPI 0.6 -
[TTHMFP]t's0
04
0.2
0
O 6.18mg03 /L-min
13 4.09 mg03/L-min
A 2.18mg03/L-min
30
TIME, mia
60
Figure 29.
Trihalomethane precursor destruction by ozone/UV:
[UV] = 0.40 W/L .
90
-------
rTTHFP\ Q6
[TTHFP]
t=o
0.4
0.2
0
O6.36 mgO,/L-min
O
G 4.18 mg03/L-min
A 2.09mg03/L-min
30
TIME, min
60
Figure 30.
Trihalomethane precursor destruction by ozone/UV:
[UV] = 0.196 W/L.
91
-------
o
M
1.2
1.0
0.8
0.6
0.4
0.2
0
0 6.36 mg03 / L MIN
Q 4.00 mg03 / L MIN
A 2.23 mg03 / L MIN
30
TIME (min)
60
Figure 31.
Trihalomethane precursor destruction by ozone/UV:
[UV] = 0.096 W/L.
92
-------
d) In the runs
I = 0.40, D = 2.18
I = 0.096, D = 4.00
I = 0.096, D = 2.23,
the primary hump appears to have been delayed from the
beginning and appears less intense than the others.
During this series of experiments, TOC was monitored as well as
THMFP. The TOC data is given in Table 9. The major point of
interest here is that comparison of Table 9 with the THMFP data
in Table 10 shows that THMFP is destroyed long before TOC is
eliminated. Thus, it appears that only specific portions of the
organic carbon macromolecules can be converted to precursor
sites, and so the model was altered to the following:
o3/uv
Precursor (P) = ^- Non-Precursor [74]
Kl
0,/UV K*
Potential Precursor Sites (M) —' 1 a^- Precursor
[75]
Non-
M _ Precursor
Integration of the resultant differential equations yields
—K t
P/PQ = e 1 (l + p/6 (l-e~6t) } [76]
where
+
p = Kl VPo
6 = KM - Kl
Note that if destruction of a potential precursor site M (KM) is
considerably faster than that of precursor (Ki) it does not take
long before e~5t is small compared to 1 and the disappearance of
P is basically exponential. If, however, precursor formation
(K+)is fast compared to destruction (Kx), some precursor is
built up before its source, M, is destroyed.
At this point the mathematical characteristics of the
problem must be considered briefly. What is desired is a model
which serves as a basis for simulating the rate equation.
Furthermore it must be a model which allows prediction of the
reaction rate, provided values of the parameters of interest
93
-------
TABLE 9. TOG DESTRUCTION IN CROSS LAKE WATER BY OZONE AND OZONE/UV
V£>
uv
Inten- Ozone
sity, Dose Rate
W/L mg/L min
2.18
0 4.00
6.00
2.23
0.096 4.00
6.36
2.09
0.196 4.18
6.36
2.18
0.40 4.09
6.18
Treatment Time, Minutes
0
8.92+0.01
9.35+0.01
8.97+0.07
8.08+0.07
8.26+0.16
7.97+0.05
8.27+0.00!
9.11+0.01
8.66+0.01
9.14+0.94
8.85+0.155
8.974+0.02
15
8.10+0.02
7.78+0.02
7.11+0.07
5.90+0.06
7.37+0.02
6.50+0.04
6.99+0.06
6.49+0.02
5.64+0.01
5.97+0.27
6.23+0.16
6.35+0.01
35
7.83+0.06
6.93+0.003
6.23+0.02
4.36+0.04
4.90+0.17
5.82+0.14
5.18+0.02
4.24+0.03
4.46+0.01
4.17+0.17
3.73+0.06
4.24+0.03
65
6.95
5.75+0.01
5.84+0.08
2.16+0.002
4.00+0.28
3.95
3.31+0.04
2.95+0.02
2.57+0.03
2.15+0.02
2.40+0.03
2.61+0.02
95
6.47+0.02
4.78+0.01
4.39+0.03
2.04+0.03
1.95+0.03
2.77+0.01
3.26+O.OOg
2.47+0.04
1.63+0.05
1.32+0.32
1.82+0.02
1.61+0.01
Pseudo-first-order
rate constant(a) (c)
-In C/C
r °i
L ^ Javg
3.39X10~3+0.03f:>)
6.54X10~3+0.030
6.70X103 +0.076
1.11X10"2 +0.004
1.38X10"2 +0.0027
1.54X10~2 +0.080
1.10X10"2 +0.074
1.34X10~2 +0.107
1.71X10"2 +0.080
1.78X10"2 +0.062
1.68X10~2 +0.126
2.01X10~2+0.060
a) Outlying values have been eliminated by the 4a rule, reference 160
b) The standard deviation should also be multiplied by the appropriate power of ten, i.e. 0.031 X 10
c) For predictive purposes, the values of K for the ozone/UV experiments (i.e. exclusive of the
ozone-only experiments) approximately fit the relationship:
K(D,I) = 2.43X10"3 + 4.29X102! + 2.96X10~4I~°'883D
-3
-------
TABLE 10. THMFP DESTRUCTION IN CROSS LAKE WATER BY OZONE/UV
ui
Inten-
sity,
W/L
.096
.196
0.40
Ozone
Dose Rate
ntg/L min
2.23
4.00
6.36
2.09
4.18
6.36
2.18
4.09
6.18
Treatment Time, Minutes
0
378 #2
348 + 33
343 + 24
352 + 20
346 + 21
290 + 18
314 + 10
282 + 10
316 + 9
5
338 + 16
305 + 12
372 + 19
397 + 23
341 + 39
232 + 12
351 + 12
358 + 3
346 + 7
10
326 + 10
306 + 12
249 + 16
391 + 20
273 + 37
119 + 22
357 + 20
260 + 19
203 + 10
20
267 + 19
215 + 17
91 + 11
326 + 14
124 + 5
32 + 17
226 + 28
111 + 19
45 + 7
30
188 + 26
128 + 3
43 + 2
207 + 15
58 + 20
45 + 17
109 + 3
35 + 14
3 + 12
45
62 + 1
84 + 5
3 + 5
78 + 8
33 + 5
1 + 19
8 + 9
60 90
33 + 9 0
33+16 0
29 + 9 0
13 + 4 0
20 + 6 0
(a) THMFP is the sum of the chlorine and bromine trihalomethanes produced by chlorination,
in yg/L. Experimental details in Appendix B.
-------
(UV intensity and ozone dose rate) are specified. Three
subordinate parameters emerge in the model in the form of rate
constants:
K, = The rate constant for precursor
destruction
p = K*M /P = The rate of formation of precursor
from site M, normalized by the amount
already present
6 = K - KI = The difference in rate constant
between site M destruction and
precursor destruction
Within any specified degree of convergence of the curve fit
and the data, there are not unique values of the parameters, K,,
p, and 6, so that if these values are iterated until they give
good reproduction of one curve, they will still not necessarily
be related to the values which fit the next curve. It is
therefore desirable to fix the form of the dependence of these
rate constants on the parameters I = UV intensity, D = ozone dose
rate, M = potential precursor site concentration, and P = pre-
cursor site concentration.
The model selected was that all reaction rates in equations
[74] and [75] were first order in substrate and had the form
-dS. .
Rate = —g^ - K.S. = k.I D S. [77]
dt 1111 L J
This is a reasonable, simple equation which has been used
successfully in the foregoing sections of this report. Since the
dependence on dose rate was obvious, the assumption of first
order in dose rate was used initially. A general fit was done
on all curves, the parameters linearized with respect to D,
curves replotted, K's adjusted to improve the fit, relinearized,
etc. This iteration was performed until the process failed to
converge further. At this point, the restriction that dose rate
be first order in KI was relaxed, and considerable improvement
was obtained. Inspection of the resulting curves indicated that
further iteration would not produce a significantly better fit
within the model used. The resulting values are
K-L = k-^D1*, where ^ = 6.38X10"2, a=0.406, b=0.727
p = k1IcDdMo/PQ, where k*M /P = .0.215, c = d = 1
KM = kMIeDf, where kM = 7.76X10~2, e = 0.325, f = 1
96
-------
where time is in minutes, dose rate in mg/L min, UV intensity
in W/L and the concentration of M and P are in the same units.
These values were used in equation [76] to produce the
curves shown in Figures 29-31, where as stated before, no
attempt has been made to reproduce the secondary humps such as
appear at t = 30 in the I = 0.196, D = 6.36 curve.
Agreement of the model with experiment is seen to be
excellent for I = 0.40 and at its worst for I = 0.196. Overall
the fit is adequate for the purpose for which it is intended
since slight variations in the matrix can cause more change in
the destruction curves than the amount by which the curve fits
miss the points in Figure 30. Of the other models tried, the
assumption that the oroginal precursor was different from that
produced by potential precursor M did not give the added flexi-
bility needed to better fit the data while still introducing
an additional variable.
It is interesting to compare the initial increase in
trihalomethane potential upon treatment with ozone or ozone/UV
with results obtained by other investigators studying different
systems. Dore, Merlet and Blanchard (98) found that treatment
of phenol with hydrogen peroxide/UV, a known hydroxyl radical
producer, or with ozone alone, increased the chloroform poten-
tial of the solution before decreasing it. Martin (179) has
reported the production of humic acid-like materials upon treat-
ment of a mixture of various phenols and carboxylic acids with
hydrogen peroxide and the enzyme peroxidase, but at this time
there has been no report of the investigation of the haloform
potential of the substances produced. In all of these systems,
there appears to be an oxygen-containing active species reacting
with aromatic ring-containing systems. Since hydroxyl radical,
for example, adds to aromatic rings rather than abstracting a
hydrogen atom, this is consistent with the m-dihydroxybenzene
precursor site theory suggested by Rook (180). In the ozone and
ozone/UV systems, precursors may even be re-synthesized from the
small molecule resulting from the degradation of larger mole-
cules. Knowledge and control of the precursor-forming part of
the ozone/UV reaction with organics could be quite important in
the optimization of precursor removal by that process, as is
obvious from inspection of Figures 29-31.
Returning to the model which was used, calculation of the
pseudo-first-order rate constants for the case I = 0.40 W/L and
D = 4.09 yields the following values of rate constants.
K = 1.22 X 10"1 min"1
p = 3.52 X 10'1 min~
KM = 2.36 X 10"1 min'1
97
-------
and the following interpretation. At t = 0, this model predicts
production of precursor occurs about three times as fast as it
is destroyed (p/Ki = pP /K,P = K+ M /l^P H 3). Destruction of
potential precursor sitis ,1-T, is however twice as fast as that
of precursor itself (which is in turn a few times faster than
the destruction of TOC), so that after a while no more precursor
can be formed. Furthermore, an increase in UV intensity
initially favors precursor formation more than it favors pre-
cursor destruction or destruction of M. The supply of M runs
out, however, whereupon the increased UV is seen to aid in the
destruction of remaining precursor.
The chemistry of trihaloraethane precursors and their
reaction with ozone and ozone/UV is obviously complicated, but
the following general scenario emerges from past work of other
researchers combined with the above observations. Humic
materials, as well as other macromolecular precursors, consist
of a diversified collection of chemical functionalities and
situations, both aromatic and aliphatic. Several of these
functionalities are known to be trihalomethane precursors.
Still others are within one or two oxidation steps of being
precursors, i.e. they are potential precursors. Evidence is
presented in the product section of this report that in hexa-
chlorobiphenyl, once a ring is broken it is degraded to a number
of products, all containing carbonyl or carboxyl groups, while
the other ring remains intact. This mechanism provides a means
of producing new oxygenated functionalities as well as reactive
double bonds which can serve as precursors as well. Balikrdshnan
and Reddy (148) have discussed the reaction of hydroxyl radicals
with benzene to produce phenol as well as a-and$-hydroxymucon-
dialdehyde. The structure of 3-hydroxymucondialdehyde, shown
below, makes it a likely precursor candidate.
Rook (15) has shown that even non-aromatic 3-diones such as
1,3-cyclohexanedione yield chloroform upon chlorination. Since
hydroxyl and many other oxy-radicals are known to add to double
bonds rather than abstract hydrogen, it is logical that as a
complicated molecule is degraded by ozonation or 03/UV treatment,
intermediates are formed which, if chlorinated, would produce
chloroform. Apparently, most of these situations are produced
early in the O^/UV reaction, and correspond to immediate con-
version of M >• P. The fact that is particularly striking,
however, is the occurrence of the secondary humps, which are
98
-------
seen in Figures 30 and 31, and even more clearly and with more
regularity .in Figure 26. A delayed accumulation of precursor
as shown in these figures, may indicate a homogeneity of some
sort in the substrate molecules, rather than a mixture of
molecular debris. If this idea were borne out by experiment,
a significant contribution would be made toward understanding
the treatment chemistry of natural water.
Summary
a) Ozone/UV destroys THMFP and TOG faster than ozone alone,
for a given ozone dose.
b) Ozone/UV appears able to entirely destroy the THMFP
of a natural water. Ozone does not.
c) Precursors in addition to those originally present may
be produced during the treatment process but are later
destroyed as the treatment continues.
d) The rate at which the THMFP of a particular water is
formed and destroyed by ozone/UV is very sensitive to
changes in the matrix which may occur over a short
time, either upon storage in the laboratory or in its
natural source.
e) The rate of destruction of THMFP in Cross Lake water by
ozone/UV could not be described by a simple kinetic
relationship of the type used for micropollutants in
previous sections. However, THMFP destruction was
described with fair adequacy by the simple model
V
Precursor 1 Mori-Precursor
o3/uv
rr
Potential Precursor 1. ^ precursor
0,/UV
Non-
M Precursor
where each reaction is assumed to be first order in substrate,
and each of the K's is described by the equation
A more accurate description of the details of the destruction
curves would require a more complicated model.
99
-------
f) THMFP is seen to be destroyed more quickly than TOC.
Mass Transfer Characteristics of the NTSU Photochemical Reactor
I. Applicability of Ozone Mass Transfer Information to the
Ozone/UV System
The most frequent criticism of early ozone work is the
failure of workers to measure the amount of ozone which was
actually consumed in a process. Most early work and even some
recent work was carried out using a glass cylinder as a reaction
vessel with a fritted glass sparger at the bottom, and either no
stirring or a magnetic stir bar for agitation. It was common to
bubble a known concentration of ozone through the solution for a
given length of time, but rarely was the concentration in the
off-gas measured, so that the actual amount of ozone reacting
was not known.
The reactor system used in this investigation was modeled
after that which has been used by Houston Research, Inc. for a
number of years. The reason for using a highly efficient
stirred-tank reactor was twofold: 1) the electrical power
necessary for the generation of ozone is a major expense, so
that the power used in stirring is more than compensated for in
ozone utilization; and 2) this type of reactor is said to be
easy to scale up (163). One important parameter in the scale-up
which must be correlatable between laboratory and full size
reactor is the ability to transfer the gaseous reactant into the
liquid. Applying the models normally used for mass transfer,
however, is complicated by several factors in the case of the
ozone/UV system. For example, it is not known whether the
majority of the ozone is photolyzed in the gas bubbles or in
the liquid. Obviously this could effect the ozone mass transfer
process, particularly if the photolysis of ozone produces a
specie which is very reactive with water. For example, in the
gas phase, photolysis of ozone produces 0 (1D) oxygen atoms which
can react exceedingly quickly with water to produce hydroxyl
radicals (82-86). For the case where 0(1D) reacts with the
bubble wall, a secondary active specie, -OH radical, would be
simultaneously produced in the liquid by the chemical reaction.
If the O^D) instead reacted with water vapor in the bubble,
hydroxyl radical would be generated in the gas phase which then
could be transferred into the liquid, react with another -OH or
some other specie in the gas, etc. If however, ozone was
photolyzed in the liquid phase, the second reaction to form
hydroxyl radical would probably occur extremely rapidly.
100
-------
/H
H 0'
\ ^ o
o'°
o o
The possibility of radical caging to form H-O- or -OOH is
obvious, and will be discussed at a later point.
In fact the identity of the active specie responsible for
the rapid destruction of substrate in an ozone/UV experiment is
not known. The three processes cited above, and others probably
contribute active specie to the solution, all with their
individual transfer characteristics. Analysis of some prelimi-
nary data collected here indicate that in our reactor, ozone
photolyzed in the gas phase is just as effective at destroying
substrate as ozone which was photolyzed in the liquid phase.
This preliminary finding indicates that the usual ozone mass
transfer considerations are not appropriate for an ozone/UV
situation. Because of the nature of this investigation and the
unavailability of resources necessary to pursue answers to the
above questions, it was decided that the ozone mass transfer
characteristics of the NTSU reactor would be evaluated to ensure
that it was performing adequately, at least by ozone mass
transfer standards.
II. Ozone Mass Transfer in Absence of UV Radiation
Since UV catalyzes the decomposition of ozone, one con-
venient way to perform consecutive mass transfer experiments is
to destroy ozone from a previous mass transfer run using the UV
lamps which were present for the ozone/UV runs. Using this
procedure it was noted that the first time a mass transfer run
was performed on a fresh batch of purified water, not only the
equilibrium level of ozone in solution, but also the calculated
value of the mass transfer coefficient was higher than on
successive runs. This effect could be explained by the pro-
duction of a substance during the UV period of the first run
which catalyzed the decomposition rate of ozone in successive
runs. Alternately, one may postulate that the organic matter
present in the water reacted with the ozone during the first
run, a situation which is known to enhance mass transfer (169,
170). In addition, the reaction might have produced compounds
or complexes which were still reactive enough to give a positive
test with iodide. The results of one such series of experiments
101
-------
is shown in Figure 32. The erratic first run followed by a
series of reproducible and regular subsequent runs made the
second possibility seem more likely. Mass transfer calculations,
described below, were performed on all runs, and in the discus-
sion which follows, the larger more erratic first runs are
labeled "first run" so that the magnitude of the effect can be
seen.
Mass transfer coefficients were calculated by the method
published by Yocum (145). His starting point was the molar
mass balance equations
G = Kla(CBe ~ CB> C?8]
- C> kc C79]
and
B _ la ,
dt V ^ Be ^B
where G = molar flow rate of the gas
Yr>r. = mole fraction of ozone in feed gas
rsr
Y = mole fraction of ozone in exit gas
K = mass transfer coefficient
J_a
C = liquid concentration of ozone in equilibrium with
gas concentration in the bubbles
CR = concentration of ozone in bulk liquid
V = volume of liqud
k = ozone autodecomposition constant
The equations resulting from Yocum's treatment are
In (l-CBn/C*) - -C^t [80]
= CB/CBe
and
kLa = r * 1 L
La S-e ] v
*
where CB = liquid ozone concentration predicted to be in
equilibrium with feed gas
102
-------
o
U)
O 7-6-79 (E-H)
3 7-12-79 IA-D)
0
10
90
100
Figure 32. Mass transfer experiments showing "first run" and subseauesnt results,
-------
C = experimental equilibrium ozone level in liquid
B = ^
P HG
L = Number of moles of liquid in reactor
H = Henry's law constant
P = Pressure
At superficial velocities (ratio of gas flow rate to
reactor cross sectional area) usually used for good mass trans-
fer into stirred-tank reactors, the volatile model compounds
chloroform, bromodichloromethane and tetrachloroethylene were
purged from the reactor at a rate which was comparable with the
reaction rate found in some of the experiments. Thus, a much
lower flow rate of 25 mL/min was chosen in order to avoid
excessive purging. This is the same ratio of gas flow to
reactor volume which was used at Houston Research in their
initial runs on trihalomethane.precursors at the beginning of
the project. Mass transfer runs were made at this gas flow rate
as well as at 200 mL/min, the flow rate used for HCB and THMFP
destruction, and at 1.0 L/min for comparison. These values are
shown in Table 11 along with the other values which were
available from the literature.
Several important points about the data in Table 11 are
noteworthy:
1) Comparison of line 1 with lines 2 and 3 indicates that
while fairly constant mass transfer coefficients are
observed at 500 rpm and 200 or 1000 mL/min gas flow,
lowering the stirring speed to 300 rpm and the gas
flow to 25 mL/min causes a decrease by a factor of two
in the mass transfer coefficient. In all of these
experiments the water had been "burned" with ozone and
UV before the experiments were performed.
2) Comparison of line 3 with line 4 shows that when the
water had not previously been burned, the mass transfer
coefficient appeared larger by a factor of five, and
comparable with the data of Yocum (lines 5 and 7) and
Prengle et_ al. (lines 8 and 9) . The two experiments in
line 4 are actually the first runs in two series, the
other consecutive runs of which are shown in line 3.
3) The data given in line 8 is the first run in the series
which produced line 10 as the second run. The numbers
in line 10 were calculated from the data presented
graphically in Figure 16 of reference 103. Since there
is no mention of UV in reference 145, it was assumed that
104
-------
the water is changed between runs. This was later
Trifi in ! PrJ-vate conversation with that author.
in other words, lines 5, 6 and 7 are also "first run"
values. The tabulated values of the mass transfer
coefficients in Table 11 are seen to break into two
groups plus the low value of run (1).
TABLE 11. MASS TRANSFER COEFFICIENTS DETERMINED FOR NTSU
REACTOR (a) COMPARED WITH LITERATURE VALUES
Gas flow
rate
ml/min
1)
2)
3)
4)
5)
6)
7)
8)
9)
10)
25
200
1000
1000
4240
4240
323
1000
1000
1000
Dose Superficial
Rate, Velocity
mg/L min cm/min
0.47
0.95-3.82
7.2
7.2
n.a.
n.a.
1.33
2,08
2.08
2.08
0.145
1.16
5.8
5.8
6.1
6.1
1.88
2.4
2.4
2.4
Stirring
Rate,
RPM
300
500
500
500
300
500
n.a.
548
720
548
g Moles
L minA x
2.69+1.53
5.26+1.71
5.30+1.66
28.1+2.3
19.2
54.5
28.6
21.3
22.9
5.30
Number of
Deter-
minations
3(f)
5(f)
6(f)
2
n.a.
n.a.
n.a.
1
1
Kf)
Refer-
ence
NTSU
NTSU
NTSU
NTSU
(b)
(e)
(b)
(c)
(c)
(d)
o ....
experimental - Appendix B.
b) Calculated by Houston Research from the data of Yocum (ref.145).
c) Given by Prengle, Hewes, and Mauk in Reference 103.
d) Calculated from the data shown in Figure 16 of Reference 103 by assuming
tQ= 130 minutes and using the three points at t=136,141, and 150 min.
CR was assured zero at t=t whereas it was actually 2.7 mg/L.
D O
e) Yocum, ref. 145, Figure 8.
f) Water which has been "burned" by ozone/UV. Results in line 4 are for
water which had not previously been "burned." See text for explanation
of "First Run" results
105
-------
la, y-moles
L-min AX
NTSU
Reference 145
Reference 103
"First Run"
Values
"Burned"
Water
28.1
5.26, 5.30
19.2, 28.6, 54.5
21.3, 22.9
5.30
The fact that the NTSU reactor behaves approximately as the
reactors described in references 103 and 145 is not surprising
since, to the extent practical, it was build according to the
design presented in those two references. The difference between
the "first run" and subsequent mass transfer coefficients,
however, is to these investigators knowledge, previously
unreported, and further study is needed to determine the reason
for this effect. The data in line 10 of Table 11 is reported to
be at 29°C rather than at 26°C which is the temperature for the
data in line 8. The change is, however, too large to be
explained by a change in the Henry's Law constant. Furthermore,
this effect is seen in NTSU experiments when no temperature
change is seen. There is one additional piece of data which
indicates that the effect is due to reaction with substances
during the first run, which are destroyed when the UV is applied
the first time. The value of the autodecomposition rate
constant, which reflects the rate at which ozone decomposes in
solution, was found to be 54.5+1.8 for the "first run" experi-
ments, while it was 15.5+4.9 for subsequent experiments using
the same water after ozone/UV treatment. Although not con-
clusive, this larger value of ko for the first run, compared
with the lower value consistently found in the subsequent runs
is consistent with the concept of "burning" the impurities out
when the UV is switched on.
Summary
a)
b)
The NTSU reactor was modeled along the lines of the
standard stirred-tank reactor design (103,145,163,168)
to the extent practical for the fabrication material
(quartz) , since "use of this design is said to
eliminate some scale-up difficulties.
Mass transfer experiments using ozone and purified
water which has never been treated with ozone appear
to give higher values of the mass transfer coefficient
than the same experiments using water which has been
"burned" with ozone/UV prior to determination of Kla.
Use of an analytical procedure for aqueous ozone which
did not respond to other oxidants might resolve this
106
-------
question, which is fundamental to ozone technology.
One. such test may be the indigo test proposed by
Hoigne1 and Bader (171). The above effect is also
seen in the published data of other investigators (103).
c) Within the framework of b) above, the mass transfer
coefficient of the NTSU reactor calculated by the
method of Yocum, does appear to be in the range
reported (103,145) for larger engineering reactors,
except at much lower flow and stir rates, as in line 1,
Table 11.
d) Above a UV intensity to ozone dose ratio, I/D, of
about 0.15 Wtmg/min)"1 no ozone or other Kl-detectable
oxidant is found in solution during an ozone/UV
experiment when either no substrate or very dilute
substrate (a few tenths of a mg/L) is present.
e) Because of d) above and the apparent ability of active
species or their precursors to transfer from the gas
into the liquid, it may be inappropriate to use ozone
mass transfer considerations to describe what occurs
in an ozone/UV experiment.
f) The autodecomposition of ozone was found to be higher
in the quartz NTSU reactor than that reported by other
investigators (103,145).
g) On a stoichiometric basis, ozone utilization for sub-
strate destruction at low concentrations (100 yg/L) is
low (Table 12) . The reasons for this low efficiency are
not known, but definitely warrant study, as a sizable
increase in efficiency would yield a very powerful
treatment tool. Along these lines it is very important
to study the mechanisms and species involved, since the
key to increased efficiency is in understanding what is
occurring in this complex system.
TABLE 12. EFFICIENCY OF CHLOROFORM DESTRUCTION BY OZONE/UV
(50% REMOVAL)
INITIAL
CHLOROFORM MOLES CHLOROFORM DESTROYED
CONCENTRATION ' MOLE OZONE DESTROYED
20 yg/L .00095
50 yg/L -0023
100 yg/L -0046
1.9 yg/L 14-5
107
-------
SECTION 6
IDENTIFICATION OF CHEMICAL BY-PRODUCTS FROM OZONE/
ULTRAVIOLET-INDUCED REACTIONS OF ORGANIC
COMPOUNDS IN WATER
EXPERIMENTAL PROCEDURES
Product runs for hexachlorobiphenyl were conducted in a
smaller reaction vessel than the one described previously for
the kinetic runs. The larger vessel was, however, used for
product runs on chloroform and lake water. No attempt was made
to determine ozone utilization except in the case of chloroform,
described below.
HCB and lake water reaction products were identified by
GC/MS after extraction and derivatization of the acidic fraction,
sometimes preceded by liquid chromatographic separation. An
electrochemical detector was used to check for easily reducible
compounds, but this investigation was cut short by the changes
in emphasis as the project progressed. Carbon dioxide was
trapped in base and titrated, and chloride ion was titrated
microcoulometrically. The details of the experimental proce-
dures are given in Appendix E.
RESULTS AND DISCUSSION
I. Chloroform, Bromodichloromethane, and Tetrachloroethylene
Product studies were not conducted on bromodichloromethane
because of limited time and the decision to deemphasize products
in favor of additional kinetic work in natural water. Because
of its lower stability and the small likelihood of finding bro-
mine and oxygen atoms on the same carbon, the products resulting
from bromodichloromethane are expected to be innocuous, and
chloroform was chosen as the representative compound.
One recurrent peak was noticed during the photolysis
kinetic studies on TCE. Attempts to extract it into ether or
pentane failed initially, but finally succeeded once the pH was
lowered to 0.8. Comparison of retention times and mass spectra
with those from an authentic sample of trichloroacetic acid
showed that they were identical. The mass spectrum is also very
similar to that of chloroform, which is expected to be the decar-
108
-------
boxylation product of the somewhat unstable acid. Further con-
firmation of the compounds identity was made from its odor.
The products resulting from ozone/UV treatment of chloro-
form were studied by treating a concentrated solution of chloro-
form (1.9g/L) with 0.57 mg/L min of ozone and 0.40 W/L of UV
radiation for one hour, during which time the off-gas from the
reactor was bubbled through sodium hydroxide solution to,trap
carbon dioxide, then through pentane to trap any purgable or-
ganics which formed. At the end of the reaction period, the
ozone was shut off and the reaction mixture purged with 75 mL/
minute of oxygen for 45 minutes to purge remaining CO2 and pro-
ducts from solution and the reactor headspace. Portions of the
aqueous reaction mixture were extracted and analyzed for halo-
genated compounds by gas chromatography with electron capture
detection, and by GC/MS. The pentane solution was analyzed
similarly, and in both cases, no compounds other than chloroform
were found. The solutions were analyzed by selected ion moni-
toring for tetrachloroethylene and hexachloroethane with nega-
tive results. The increase in carbonate in the trap was mea-
sured and a portion of the reaction mixture titrated for hydro-
gen ion increase. Another portion of the reaction mixture was
derivatized and analyzed for formic acid as the benzyl ester,
as described in Appendix E. The results are shown in Table 13.
TABLE 13. PRODUCTS RESULTING FROM OZONE/UV TREATMENT
OF CHLOROFORM
Concentration Ratio to
Substance m moles Chloroform
Chloroform (destroyed) 7.32 + 0.73
H+ 23.6 + 0.4 3.22 + 0.38
Cl~ 19.4 + 0.2 2.65 + 0.29
C02 5.1 + 0.2 0.70 ,+ 0.10
Halogenated Organics n.d.
Formic Acid n.d.
Ozone Destroyed 0.713 0.097
109
-------
It can be seen from the last column that the reaction
CHC13 ^ 3H+ + 3C1" + C02
appears to have occurred. Within experimental error the pro-
ducts are those expected from the equation (Cl~ lacks 3% of ful-
filling the stoichiometry while C02 lacks 20%, probably due to
the method of collection), thus it is seen that the products are
inorganic and innocuous, organic halogen being mineralized to
halide ion. Several mechanisms can be written which would give
this stoichiometry, but these are the products which are ex-
pected whenever one of two things happens: 1) an oxygen atom in
any form is attached to the carbon, whereupon an acid chloride
situation develops, which is unstable enough in water that the
acid chloride itself (HCOC1, C1C02H, COC12, etc.) probably never
forms, but passes on to the hydrolysis products shown in the
above equation; or 2) the hydrogen atom is abstracted by a free
radical in the presence of water (the autooxidation product of
chloroform is phosgene). In either case, once the integrity of
the chloroform molecule is destroyed it probably collapses com-
pletely to products, resulting in the very rapid reaction rate
observed and the lack of intermediate products. Judging from
the last entry in Table 13, this appears to be a chain reaction.
It is not presently known what species carries the chain, but
possible coupling products between CHj_Clj species were specifi-
cally looked for in the chromatograms of both reaction mixture
and pentane trap and not found. The above considerations prob-
ably hold for most halogenated compounds, consistent with the
claim by Westgate Research Corporation for example, that ^their
jszone/UV sy_stem is_part.icularly eff_ective_ in removing chlori-
nated "compounds (115).~ " """—— —
II. Hexachlorobiphenyl
As an aid to understanding and predicting what types of
products will be produced upon ozone/UV treatment of complex
halogenated molecules, the degradation of 2, 2', 4, 4', 6, 6'-
hexachlorobiphenyl was studied rather extensively. Although the
concentration of PCS in water and wastewater is usually at the
yg/L level, the experiments in this work were conducted at
higher concentration (suspension) in order to obtain a larger
quantity of products for identification. The small reactor as
described in the experimental section was used for all the
experiments. In the case of photolysis, nitrogen or oxygen gas
at the same flow rate was used. At the end of each reaction
time, the whole solution was extracted according to the proce-
dures described in Appendix E. Products were identified by
using gas chromatography combined with mass spectrometry. Chro-
matograms and mass spectra are collected in Appendix E, section
II, but are referred to from the text.
110
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Photolysis Products of HCB—
The photolysis of HCB in nonaqueous solution has been
studied by Safe and Hutzinger (172) with 310 nm radiation. The
authors found that chlorine is removed rather rapidly from HCB
yielding di-, tri-, tetra- and penta-chlorobiphenyls. The
reconstructed GC/MS scans of HCB photolysis products found in
the neutral fraction in this work are shown in Figures E-3 and
E-4. No acidic products were found. In the neutral fraction of
the solution after four hours radiation, dechlorinated product
pentachlorobiphenyl was found and heptachlorobiphenyl was also
identified. Safe and Hutzinger found other dechlorinated pro-
ducts besides pentachlorobiphenyl and some other compounds
caused by rearrangement or condensation. However, in this work
only products resulting from dechlorination were found. The
product heptachlorobiphenyl is believed to be produced from the
attack of HCB by chlorine radicals. It should be noted that the
solution used is a suspension with high concentration of HCB and
the heptachlorobiphenyl is an unlikely product in dilute solu-
tion. Figure 33 shows the production of chloride ion during the
photolysis process as measured by microcoulometry. The yield
appears to be approaching 1 mole of chloride ion per initial
mole of HCB.
When the HCB molecule is excited by UV radiation, the C-C1
bonding is apparently cleaved. According to Ruzo et al (173),
this process proceeds through the triplet state and is driven
by the need to remove strain in the excited molecule in which
the planar configuration is most stable. The formation of
heptachlorobiphenyl suggest that free chloride radicals are
formed (path A). However, it is possible that the pentachloro-
biphenyl product is the result of a concerted process in which
solvent water attacks the excited HCB molecule (path B),
Figure 34.
The mass spectra of HCB, penta chlorobiphenyl and hepta-
chlorobiphenyl are shown in Figures E-5, E-6 and E-7. From
Figures E-3 and E-4, and from the rate data reported before/it
is apparent that oxygen does not change the rate or mechanism
of the photolysis of HCB in aqueous solution. It has been
reported that the presence of oxygen in nonaqueous solution
retards the photolysis of PCBs, presumably due to triplet
quenching by oxygen.
Ozonation and Photolytic Ozonation Products—
Yokayama and co-workers (174) ozonated a mixture of PCBs
in aqueous solution for 24 hours to obtain a variety of small
degradation products which were rationalized in terms of a ring
opening process. Chlorine elimination was found as a secondary
process, and it was noted that PCBs with several chlorine atoms
produced fewer products. No detailed study of the ozonation
products of HCB has been conducted in this work; however, in one
ozonation run, traces of pentachlorobiphenyl were found along
111
-------
0.3
[cr]
(M)
0.2
O.I
0
TIME, hr
O
Figure 33. Production of chloride from the photolysis of HCB in water.
-------
with a compound with M = 372 with a six chlorine cluster. The
compound was not identified. No chloride production during the
reaction period of four hours was found.
Photolytic ozonation of HCB was conducted in the small
reactor and the products identified by either GC or GC/MS. The
reactor is shown and extraction and derivatization schemes
described in Appendix E. In order to study the pathways of the
reactions, the concentrations of the major products at different
reaction times were determined. The proposed reaction scheme
H
Cl-
OH
(CHLORINE IS SHOWN BY A LINE — )
FOR SIMPLICITY
Figure 34. Hexachlorobiphenyl photolysis pathway according to
Ruzo (173) (chlorine is shown by a line - for simplicity).
113
-------
for ozone/UV degradation of HCB is shown in Figure 35. The
scheme utilized five types of reactions which are not uncommon
in ozone (176) and photochemistry (175). Double bonds are
cleaved in aqueous solution (C in Figure 35) to yield acids and
carbonyl compounds. Decarboxylation (D in Figure 35) and re-
placement of halogen by hydroxyl (R) are known photochemical
processes (175). Oxidation of an aldehyde to an acid and hy-
drolysis of an acid chloride are the remaining two reactions
which are employed in the scheme. This scheme is not intended
to rationalize all of the products which were found, but is
meant to be representative of the type of stepwise degradation
which is occurring.
The reaction scheme proposes that one ring is first rup-
tured and then degraded in a series of rapid steps before the
second ring is attacked (177). The repetitive process which
destroys these compounds is, a) double bond cleavage resulting
in an aldehyde which, b) oxidizes to a carboxylic acid, which
c) decarboxylates. Further cleavage may occur when a molecule
contains additional double bonds, so that organic carbon in
converted to carbon dioxide via either intermediate carboxylic
acids or the terminal compound in the ozone process, oxalic
acid. Since most of the products are carboxylic acids, pH pre-
sumably plays an important role in the degradation of products.
At elevated pH, the acids are dissociated and more reactive
toward electrophilic reagents; however, pH effects in the pro-
duction of products were not studied in this work.
The compound 2-chloro-3-(2,4,6-trichlorophenyl)-maleic
acid, (CTMA) is a major intermediate which is produced rather
quickly and reaches a maximum at about 3 hours. This compound
is believed to be the first stable ring-opening product of HCB.
Upon ring opening the HCB molecule apparently quickly oxidizes
by the loss of two carbon atoms in the form of oxalic acid,
which also is produced at a high rate as shown in Figure 36.
2,4,6-trichlorobenzoic acid forms steadily as the reaction pro-
ceeds, and is relatively stable compared to other products. It
is significant to note, however, under UV radiation decarboxyl-
ation (175) occurs and the product 2,4,6-trichlorobenzene is
observed.
After ring-opening of HCB further oxidation seems to pro-
ceed rapidly on the open chain, leaving the second ring undis-
turbed. The formation of chloride in the photolytic ozonation
system occurs -steadily as shown in Figure 37. The percentage
of chloride relative to the theoretical value is slightly high,
possibly due to the interference of products in the microcoulo-
metric titration reaction. Nonetheless, the steady formation of
chloride ion suggests that chlorine on HCB may be attacked to
initiate ring rupture and subsequent oxidation reactions.
Although peroxides are assumed intermediates from the oxidation
of organics, high pressure liquid chromatography with an elec-
114
-------
CI*
CI2
r-t W
-CH>H
N ( >—OH
0 0
O
OH
01 0
OH
OH
CH,
CH2CI
RING FRAGMENTATION
CU
CH3
extending from the benzene ring).
115
-------
9
o
o
X
O
0
2
hr
Figure 36.
Formation of oxalic acid during photolytic ozonation
of HCB in water.
116
-------
reaction time » hours
Figure 37. Chlorine balance during ozone/UV treatment of HCB.
-------
trochemical detector failed to detect easily reducible peroxides
as HCB oxidation products. Presumably these intermediates, if
formed, are very reactive and are destroyed very quickly under
the severe photolytic ozonation conditions. Appendix E contains
typical reconstructed gas chromatograms (Figure E-8) and GC/FID
chromatogram
products. All compounds from CTMA and beyond in Figure 35 were
seen in the GC/MS data, with the exception of trichlorophenyl-
acetaldehyde.
The major compounds are listed in Table E-l, the amount of
products produced as a function of time is shown in Figure E-10,
and the mass spectra and the fragmentation patterns on which the
identifications were based are shown in Figures E-ll to Figure
E-19.
III. Natural Water: Caddo Lake, Texas
In considering the ozone/UV treatment of natural water to
remove micropollutants and trihalomethane formation potential,
it is o:f importance to know what kinds of products are formed
from the ozone/UV reaction with the matrix itself. Caddo Lake
water has been ozonated with and without UV, products extracted,
and studied by GC/MS. The experimental details, reconstructed
mass chromatograms, and lists of compounds identified are pre-
sented in Appendix E and summarized below.
The only peaks which have been identified are those in the
ozone and ozone/UV runs which are not present in the control.
Methods of confirmation of peak identity include comparison
with spectra collected in Registry of Mass Spectral Data (164)
or published in the literature, and computer matching through
the Chemical Information System (166-) . Those spectra which
could not be found in the literature have been rationalized
according to fragmentation patterns and compared with actual
samples whenever possible. The scheme shown in Figure E-2 for
concentration of HCB products is also the extraction scheme used
to concentrate products from the Caddo Lake water ozonations,
neutrals and basic compounds being extracted at high pH and
acids at the lower pH. Glass capillary mass chromatograms were
obtained by splitless injection onto a 30 m. SP-2100 WCOT column
in a Finnigan 3200 GC/MS system, and are shown in Figures E-18
through E-23. Peaks which occur in the ozonation runs but not
in the control are labeled by letter and the compounds are
listed in Table- E-2. In many cases, the compounds have not been
identified but several peaks are seen to exhibit the same frag-
mentation pattern. In these cases the fragmentation patterns
are listed in the notes to Table E-2.
The most striking feature of the mass chromatograms shown
in Figures E-18 through E-23 is that treatment of Caddo water
with ozone greatly increases the number of lower molecular
118
-------
weight (i.e. chromatographable) compounds, as can be seen by
comparing Figure E-18 with Figure E-20 and Figure E-21 with
Figure E-23. Ozone/UV treatment, however, produces much fewer
such compounds, many of which are thought to be innocuous at low
levels (e.g. aliphatic hydrocarbons) and most of which are at
much lower levels than, with ozone alone. (Compare Figures E-19
and E-28 with E-18, E-20, E-21, and E-23.) Quantitation has not
been attempted but peak f is the internal standard, trichloro-
benzene, the height of which corresponds to 1 yg/L in the orig-
inal water, assuming 100% extraction efficiency.
In order to avoid altering the nature of the products, the
remaining ozone residual was not quenched by the addition of
some reducing agent. Peak C is thought to be related to ethyl
ether peroxide and has the empirical formula C4HgC>2. Striking
is the number of nitrogen compounds found in the chromatograms,
but more expected is the abundance of organic acids produced in
the experiments using ozone without UV. Hydrocarbons in the
03/UV neutral fraction are though to originate by photolysis of
carboxylic acids (175) ,
hv
RCO0H
RH
CO-
or by reaction with an active free radical. Seivers (60) has
reported their production in a non-UV ozonation system. Absence
of organic acids in the 03/UV acid fraction is conspicuous, and
is consistent with the findings of other investigators (103).
Of interest from a health effects viewpoint is compound n.
Fetizon and Audier (14) reported the mass spectra of I and II
and discussed the possibility of their interconversion through
the epoxide III. The spectra obtained in our experiment was
intermediate between the published spectra of I and II, making
the presence of the epoxide a possibility. There is, however,
no direct evidence for the presence of the epoxide.
II
III
Summary
It can be seen from the ozonation products of Caddo Lake
water that the number of chromatographable products, both from
ozone and from ozone/UV, is larger than that found for the con-
trol, indicating breakdown of larger natural organics into
smaller molecules. Ozone/UV treatment, however, while showing
considerably greater destruction of the organic matter (TOC)
119
-------
present, gives considerably fewer products than ozonation alone.
Seen from the work on hexachlorobiphenyl is the fact that once
an aromatic ring is ruptured, it appears to continue to degrade
more rapidly than scission of the other ring occurs. Further-
more, with HCB, as in the case with chloroform, chlorine atoms
are eliminated from the degrading portion of the molecule very
early in the process, as the first oxygen atoms are introduced.
Because of the decrease in TOG in the case of chloroform, it is
seen that organic compounds actually are converted to carbon
dioxide and organohalogen to halides by ozone/UV, as opposed to
the situation for ozone alone, where degradation usually stops
at some organic compound such as oxalic acid. This difference
is reinforced by comparison of the ozone and ozone/UV products
from natural waters, referred to above. One strength of the
ozone/UV process over ozone alone seems to be the ability of the
ozone/UV process to decarboxylate organic acids.
Comparison of the effect of ozone and ozone/UV on aqueous
chloroform is particularly instructive since many of the early
investigations of nonaqueous ozone chemistry were performed
using chloroform as a solvent. The ozone/UV process is seen
from the present work on chloroform to involve a chain reaction,
at least in the case of some substrates. It was observed that
several chloroform molecules are destroyed by photolysis of one
ozone molecule at high substrate concentrations, while at very
low substrate concentrations the ratio of chloroform to ozone
molecules destroyed went as low as 10~3 for 20 yg chloroform per
liter. This observation is in agreement with the chain reaction
hypothesis in that the ratio of the rates of the chain termina-
tion reactions to propogation reactions is very large when the
chain carrier finds it difficult to locate substrate. Similarly,
when a natural matrix is added, additional possibilities for
chain termination and other reactions are introduced.
120
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SECTION 7
PROCESS ENGINEERING ANALYSIS AND SIZING AND
COSTING OF FULL SCALE UNITS*
A consulting engineering firm was hired to design a full
scale process for removal of chloroform, bromodichloromethane,
and tetrachloroethylene from lake water in one set of cases, and
to remove trihalomethane precursors from lake water in a second
set of cases, utilizing the kinetic data collected in this
project and presented in preceding sections. This design would
be utilized by the firm to estimate treatment costs for the
process. As test cases, 1, 10, and 50 million gallon per day
capacities were selected, and facilities were designed to remove
90% of the undesirable substances.
This section contains the final report submitted by the
subcontracted engineering firm, Houston Research, Inc., selected
as representative of the best currently available technology,
due to their experience in the field of ozone/UV treatment. The
critical assumption which has been made in performing the
following calculations is that of how much the efficiency of the
process can be increased in going to full scale equipment.
Houston Research, based on their previous experience, has
estimated that an ozone utilization efficiency of 35 to 50% may
be expected from the full scale unit, whereas maximum effi-
ciencies of 17% in the case of trihalomethane precursors and 13%
in the case of the micropollutants were calculated from the
experiment, as based on reaction stoichiometries assumed by
Houston Research. Whether this difference is due to the low
concentration of substrate, some characteristic of the laboratory
scale reactor, or some other factor is presently unknown.
However, efficiency calculations performed using data similar to
that shown in Figure 32 but with a model substrate present,
indicate that all ozone photolyzed, whether in the liquid or the
bubbles, was utilized to the same extent in the destruction of
model compound. These results indicate that a more likely cause
of the low efficiency was the competition of side reactions with
useful ones. While examination of that question was beyond the
*This section of the report was prepared by H. William Prengle,
Jr., Arthur E. Nail, and Dilip S. Joshi of Houston Research,
Inc., Houston, Texas.
121
-------
scope of this investigation, the study of the mechanism and
active species of the ozone/UV process would be of importance
in the process optimization. A pilot study would be needed
before the assumed increase in efficiency could actually be
demonstrated.
The following list defines the symbols and nomenclature
used in this section.
122
-------
SYMBOLS & NOMENCLATURE
Chemical Species
A .... TOG or other organic components
E Ozone
C Chloroform (CF)
D Bromodichloromethane (BCM)
E Tetrachloroethylene (TCE)
M Trihalomethane Precursors (THMP)
S Substrate of interest; or total
of all organic species
Upper Case
BC • *• Cumulative ozone supplied
CB/CB Concentration of ozone; Henry's
law concentration of ozone,
respectively
D Ozone dose rate
D.J.,DT Impeller diameter; tank or
reactor diameter, respectively
E , E _ UV energy; UV energy per gram of
S removed
H Henry's law constant
H , E, , H£, Height of impeller above bottom
of reactor; height of liquid,
height of sparger, respectively
I UV-intensity
K, , K _, K^, K_, K,, Component chemical reaction rate
A C +D E M constants
K,K ,K ,K Reaction rate constant; for 03-
0 u u only; for UV only; preliminary
UV rate constant, respectively
K ,K Overall mass transfer coef-
L 9 ficients, liquid side, gas
side, respectively
K* UV-energy transfer coefficient
LV Path length
N,N ,N ,N Rotation rate; gas flow number;
9 R B Reynolds number; 0, mass
transfer number, respectively
P,P Pressure; power input
Q Gas flow rate
R,r Reactor radius; bulb to reactor
distance, respectively
V Reactor volume
Xx* Ozone liquid phase mole fraction;
B' B Henry's law value, respectively
123
-------
Lower Case
a Interfacial area for mass trans-
fer; liquid phase attenuation
coefficient, respectively
b,b,b' Turbine blade width; apparent 03
a requirement; stoichiometric
03 requirement, respectively
f (N ) Gas flow dispersion function
g,gg UV irradiation geometric factor;
c dimensional constant (fm Ibm/
Ibf sec2); respectively
h Turbine blade height
k ,k Individual mass transfer coef-
g ficients, liquid and gas,
respectively
k a,k a Mass transfer coefficients
. ^ including interfacial area
mB,m Ozone supply rate; volatile
component desorption rate,
respectively
n UV-intensity exponent
p^ Vapor pressure
r Correlation coefficient
t Time variable
u . . . Superficial gas velocity
w, Reactor quadrant baffle width
wn,w Nominal UV bulb wattage; UV-
wattage, respectively
y Vapor phase mole fraction of
ozone
Greek
a,a , a 0.,-utilization fraction; for 03
only; intrinsic value
Y Liquid phase activity coefficient
T] Preliminary analysis optimi-
zation parameter
y Liquid phase viscosity
p Liquid phase density
T Wall transmission factor
Operators
< > Denotes average
Sub o Denotes initial value
Super o Denotes standard state values
"Dot" above symbol denotes a
time derivative
124
-------
A. PROCESS ENGINEERING ANALYSIS BASIS
The Ozone/UV process for photo-oxidation of refractory and
toxic inorganic and organic species was developed in the early
1970's. The first quantitative reaction rate data were pre-
sented in 1973 [1], and followed by extensive data on oxidation
of organic compounds in 1975 [2] and 1976 [3,4]. A substantial
body of research has been completed on a broad range of inorganic
compounds: metal cyanide complexes; organic and amino acids;
alcohols; insecticides; organic nitrogen, phosphorous, and sul-
fur compounds; chlorinated organics; inorganic anions and
cations; and numerous composite industrial mixtures. Three
major areas of application exist for treatment of: a) indus-
trial waste water, b) municipal waste water and c) raw source
water. To date, three patents have been granted [5,6,7].
And now detailed research, presented in this report, has
been carried out on application of the process to removal of
micropollutants and trihalomethane precursors from raw water.
The objective of this part of the report is to carry out a
reaction engineering analysis of the experimental data, adding
the necessary important aspects of reactor dynamics, in order
to produce,
1) - a process configuration, and
2) - sizing and costing of the equipment for certain full-
scale units.
The work presented in the first part of the report is not
sufficient, of itself, to permit design of the reaction equip-
ment. Other aspects, such as mass and energy transfer and
optimization must be added in order to do the full-scale designs.
Prengle has developed the methods, used herein, from hundreds of
Houston Research, Inc. 03/UV data points and basic chemical
engineering theory and practice.
At the outset, some comments are necessary concerning the
reaction rate equations used. Peyton, Glaze and co-workers in
the first part of this report have used the forms, e.g. for the
tetrachloroethylene data,
o _ K IDS (la)
-fa - A
,iaols S. , £ min ) (watts, (^3 )f (Sols*
l£ sec '' lsec,mg03,watts/A ' S- * min
integrating to,
£n SQ/S = Ku I D t
125
-------
and gives a hybrid non-consistent set of dimensions for the over-
all reaction rate constant, Ku, (5L, min/sec, mgO3, watts/*). An
alternate form has been used in the reaction engineering analysis
that follows:
K I l-^J (-Ay) <2a)
u V s
mg O_
mg S > .watts. ,_ 3_>
watt/*' ' ( I '' * min '
S°, a standard state concentration (taken as 1.0 mg/*) is
inserted to make the concentration variable dimensionless.
The equation integrates to,
*n(S /S) = K I(m_/V)t = K IB (2b)
O U B U C
Since the sequence of reactions taking place by photolysis and
photo-oxidation are driven by UV-energy input, the ozonalysis
being almost negligible, the dimensions of Ku[(mgS removed/mg03,
watts/*)] = [(mgS removed/mgO3, (joules/sec *)], is a dimen-
sionally consistent concept.
B. ENGINEERING ANALYSIS OF THE REACTION DATA
Each set of experimental data obtained was analyzed to
determine certain quantities necessary to carry out the process
engineering calculations and sizing of the full scale units.
The characteristics of the lake water used in the experi-
ments are shown in Table 14.
TABLE 14. CHARACTERISTICS OF THE LAKE WATER
Characteristic Lewisville Lake (A) Cross Lake (B)
pH
SS (mg/*)
TDS (mg/*)
COD (mg/*)
TOD (mg/*)
THM Precursor (yg/*)
COD/TOG
8.1
<10
130
10.0
4.3
-
2.33
7.6
<10
100
34.6
8.7
<330>
3.98
An important parameter missing from the above list is a measure
of the "total oxidizables, TO," or the total oxygen demand, TOD.
Such information is very important in determining the total
0-,-utilization, since Houston Research, Inc. experiments have
126
-------
amply demonstrated that inorganic species — certain metals,
cations, anions, etc. are also oxidized by 03 and 0,/UV, pro-
ducing an additional ozone demand. However, at present a
satisfactory TO test procedure is not available. The data sets
are analyzed in Table 15.
TABLE 15. ANALYSIS OF DATA SETS
Initial Concentration
System (mg/£) Table
1)
2)
3)
4)
5)
6)
Tetrachloroethylene ,
TCE in purified water
TCE in Lake Water (A)
Chloroform,
CF in purified water
CF in Lake Water (a)
Bromodichlorome thane ,
BCM in Lake Water (a)
Trihalomethane Formation
Potential (5 days, 26°C)
0
0
0
0
0
<0
.100
.100
.100
.100
.100
.330>
II-l
II-2
II-3
II-4
II-5
II-7
In order to carry out the engineering analysis of the data
the following functions were used as "screening" parameters.
1) The cumulative 07 (B) supplied,
Bc = <-)t, (mg03 supplied/ i) (3)
2) The apparent 03 requirement,
B
b (based on S) = (T|) , (mg03/mgS removed) (4a)
a at> J
b (based on TOO = (j^) , (mg03/mg TOC removed) (4b)
a
3) The fraction of 0 utilized,
3
(b'ASj ^ b, _ stoichiometric 03 (5a)
a (on S) = Bc ' required
.COD. A (TOC) (5b)
a (on TOO = (rfoc'o ~g
c
127
-------
4) The reaction rate constant,
Ku(on S) = £n(So/S) , mg S removed . (6a)
I (B _/B°) mgO,, joules/sec l'
C j
removed (6b)
K (on TOC) - i
Von TOC) = I(BC/BU) ' {mg03, joules/sec
These K -rate constants used for screening assume first order in
I and Bu; therefore, for a given set of runs, a drift in the K -
value will be an indicator of whether or not first order is u
adequate to account for the results. In addition, the question
arises: How constant is constant? Some take the view that if
rate constants fall within "an order of magnitude," they are
constant, and there are literature data for gas phase reaction
which only satisfy such a criterion. Here, at the outset, we
would consider values constant if they are within a factor of
2 to 3.
5) Energy input per unit volume,
_ _ /Itx ,watt hrs. ._ .
E = (60> ' { - 1 - J (7a)
6) Energy input per gram of material removed,
Ev
E = ( - = — ) (watt hrs/g. S removed) (8a)
VS 10"3AS
E
EMTn<- = ( — rr^ - ) ' (watt hrs/g. TOC removed) (8b)
ViUC 10
7) An approximate optimization parameter can be defined and
calculated from the data based on the following concept.
Best process economics for a given situation would dictate
that the size of reaction equipment, amount of ozone, and
amount of energy input be minimized. That is, maximum S or
TOC removal, per unit of '03 input, per unit of energy input,
per unit residence time. The requirements can be incor-
porated in the parameter, r\ , which is used only for screening
purposes,
= t AS rcg S removed
'
B E t ' mg0-., (watt hrs/£)min
G v j
128
-------
Analysis of the individual data sets follows.
1. Tetrachloroethylene Destruction in Purified Water
Example; I = 0.375 watts/A, (m/V) = 1.33(mg 00/min) ,
13 J
t = 0.917 min
SQ = 0.100 mg/SL, AS = 0.095, £n SQ/S = 3.00
B = (m /V)t = (1.33) (0.917) =1.22 mgO_ supplied/A
C JD j
b = (B /AS) = 1.22/0.095 = 12.8 (mgO-./mg AS)
cL C j
&n So/S _ 3.00 _ fi ,., ,mg S removed
l
_ _
u T/_ /Do. (0.375) (1.22/1.00) mgO^ watts/A
/ & I 3 I
a = k^S = (1.158)(0.095) = Qi(J902 (fraction Q3 utilized)
15 _L • ^ ^
C
E = (It/60) = (°'37?I (°-917) = 0.00573 watt hrs/£
V 60
E _=
Ev 0.00573 _ cn _,watt_hrs.
— -
_ - •=. — T
VA AS(10"3) 0.059(10-3)
AS . 0.095
_ = _
(1.22) (0.00573) (0.917) • mg0 (watt hrs
J j x/
Values of these parameters for all conditions of TCE in purified
water are summarized in Table 16.
2. Tetrachloroethylene Destruction in Lake Water (A)
Example :
I = 0.375, (mB/V) = 1.33, t = 11.1 min, SQ = 0.100,
AS = 0.090, TOC0 = 4.3, (TOC/TOCQ) = 0.81,
ATOC = 0.82
129
-------
TABLE 16. SUMMARY DATA ANALYSIS, TETRACHLOROETHYLENE
DESTRUCTION IN PURIFIED WATER
I 0 t
(watts/A ) (mg03/min) (min)
0.091 1.33
0.775
0.420
0.200
0.090
0.189 1.33
w 0.775
o
0.460
0.230
0.110
0.375 1.33
0.775
0.450
0.260
0.090
9.83
7.83
16.7
-
33.0
1.53
2.58
5.33
9.33
10.1
0.917
1.57
2.23
3.17
6.60
B
C
(mg03/£)
13.1
6.07
7.01
-
2.97
2.03
1.99
2.45
2.14
1.11
1.22
1.22
1.00
0.824
0.594
b
a
(mgOVragAS)
136
63.2
73.0
-
3094
21.1
20.7
25.5
22.3
11.6
12.0
12.7
10.4
8.58
6.19
K ni ",.— fl
U a VS '
(watts/A)"1 (watt hrs/gAS)
2.52
5.43
4.70
-
11.1
7.82
7.98
6.48
7.41
14.3
6.56
6.56
8.00
9.71
13.5
0.0085
0.0183
0.0158
-
0.0371
0.055
0.056
0.045
0.052
0.100
0.0902
0.091
0.111
0.135
0.187
155
124
264
-
521
50.2
84.7
175
307
332
60.3
102
146
207
431
0.0495
0.168
0.0323
-
0.0194
6.42
2.30
0.438
0.164
0.268
14.8
5.11
3.09
1.86
0.593
Constants for this Set of Data: b' = 1.158, SQ = 0.100 mg/&, &rt(S /S) = 3.00;
As = 0.095 mg/A.
-------
Bc = (1.33) (11.1) = 14.76 (mg03/£min)
ba(on S) = (Bc/AS) = (14.76/0.090) = 164
ba(on TOC) = (Bc/ATOC) = 14.76/(0.19) (4.3) = 18.07
0 Y"\ C! / C O *D r\ o
K (on S) = n 0/~' = 2-JOJ _ o.416
U I(BC/B°) (0.375) (14.76)
= (1.158) (0.090) = Q
a(on TOC) = >76 = 0.129
E = It = (0.375) (11.1) _
EV - 6Q - go - - 0.0694
E = Ev = 0.0694 (1000) = 771
ASdO'3)
„ = EV _ _ (0.0694)103 69.4 . ,
VTOC " ATOC(10-3) " °'82 " °-82 "
A TOC 0.82
n
_
TOC B E ,t " (14.76) (0.0694) (11.1)
C V
Values of these parameters for all conditions are summarized in
Table 17.
3. Chloroform Destruction in Purified Water
Example ; (CFM in purified water)
I = 0.184, (mB/V) = 0.733, t = 12.0, SQ = 0.100
AS = 0.090, Jin SQ/S = 2.303
B = 0.733(12.0) = 8.80
c
b = (8.80/0.090) = 97.8
a.
131
-------
TABLE 17. SUMMARY DATA ANALYSIS, TCE DESTRUCTION IN LAKE WATER (A)
- Based on S -
I
/"B ,
( V ;
) t
(watts/ (mgO3/ (min)
&) min)
0.084
0.189
0.375
1.33
0.775
0.510
0.295
0.11
1.33
0.775
0.480
0.230
0.10
1.33
0.775
0.420
0.260
0.096
Constants for
20.3
29.5
40.4
52.9
72.3
14.6
22.1
29.3
41.0
51.7
11.1
16.2
23.0
28.2
37.3
this Data
Bc
,TOC,
TOC0
(mgoy
27.0
22.9
20.6
15.6
8.0
19.4
17.1
14.1
9.4
5.2
14.76
12.5
9.7
7.3
3.6
Set:
0.92
0.91
0.91
0.93
-
0.87
0.86
0.87
0.91
0.94
0.81
0.81
0.83
0.84
0.92
b1 =
b
a
(mgCX/
mgAA)
300
254
229
173
89
216
190
157
104
58
164
139
108
81
40
1.158, S
K
u
a
EVS
ba
Based on TOC -
K
u
a
EVTOC
n
(watts/
1.02
1.20
1.33
1.76
3.12
0.628
0.713
0.863
1.30
2.35
0.416
0.492
0.633
0.842
1.56
<1.22>
= 0.100
0.00386
0.00456
0.00505
0.00665
0.0131
0.00855
0.00610
0.00740
0.0111
0.0202
0.00706
0.00805
0.0108
0.0142
0.0292
mgA, S
316
459
629
823
1124
511
773
1026
1436
1811
771
1165
1597
1958
2590
78.5
59.2
53.2
51.8
-
34.7
28.4
25.2
24.3
20.2
18.1
15.3
13.3
10.6
10.5
= 0.090 mgA
0.0367
0.0490
0.0545
0.0554
-
0.0380
0.0467
0.0523
0.0531
0.0630
0.0382
0.0450
0.0512
0.0637
0.0618
<0.0506>
0.0297
0.0394
0.0438
0.0450
-
0.0671
0.0820
0.0924
0.0959
0.1156
0.129
0.152
0.176
0.220
0.223
, to S /S = 2.
34.6
50.4
69.0
90.4
-
56.1
84.9
113
158
199
84.6
124
175
215
284
303
0.0221
0.0139
0.0082
0.0049
-
0.0429
0.0229
0.0147
0.0078
0.0059
0.0721
0.0400
0.0256
0.0262
0.0110
For TOC:
TOC = 4.3 mgA,
2.33
-------
An S /S 2.303
K = - - — = - = 1 42
U I(B/B°) (0.184) (8.80)
C
_ b'
a-
b'AS _ (1.004) (0.090) _
ng— - - (g-- - - 0.0103
c
= (0.1S4M12.0) .
EVA= -3 - °-°368 .3 - 409
AS (10 ) (0.090 x 10 )
AS 0.090
_ _
B E t (8.80) (0.0368) (12.0)
C V
Values of these parameters for all conditions of chloroform in
purified water are summarized in Table 18.
4. Chloroform Destruction in Lake Water (A)
Example ;
I = 0.375, (mD/V) = 1.33, t = 171 min, S = 0.100,
D "J
AS = 0.090, (fcn S /S) = 2.303, TOCQ = 413,
(TOC/TOC ) = 0.037; ATOC =4.14
B = (m^/VJt = 1.33(171) = 227
C^ -D
b (on S) = (B /AS) = (227/0.090) = 2522
3, C
b (on TOO = (B /ATOC) = (227/4.14) =54.8
3. C
£n S /S 2.303
K (on S) = - °— = - = 0.0271
U I(B /B°) (0.375) (227/1)
C
Jin (TOC/TOC) 3.297 nn,87
V°n TOC) ' o - ' 0.375(227/1) - °-0387
133
-------
TABLE 18. SUMMARY DATA ANALYSIS, CHLOROFORM DESTRUCTION
OJ
J-Vf , r l«L£U~C..J-Cd-' . .MLa-l Tj R . . — - - —
m
I ( v )
(watts/5,) (mgO3/
min)
0.084 0.716
0.545
0.443
. 0.330
0.184 0.733
0.545
0.483
0.458
0.318
0.417 0.625
0.477
0.364
0.238
t
(min)
15.4
23.1
31.5
49.2
12.0
1818
22.5
24.3
42.1
12.5
18.4
27.7
52.1
B
c
(mgO.,/&)
11.0
12.6
14.Q
16.2
8.80
10.3
10.9
11.2
13.4
7.69
8.78
10.1
12.4
ba
(mgO,/
mgAAJ
122
140
156
180
97.8
114
121
124
149
85.4
97.6
112
138
K
u
a
(watts/ A) ~l
2.49
2.18
1.96
1.69
1.43
1.22
1.15
1.12
0.934
0.718
0.629
0.547
0.445
0.00910
0.00795
0.00715
0.00600
0.0103
0.00880
0.00830
0.00810
0.00675
0.0118
0.0103
0.00895
0.00725
EVS
n
(watt hrs/ (X 10 3)
gAA)
240
359
490
765
409
640
767
829
1435
950
1421
2140
4020
0.0246
0.00955
0.00463
0.00163
0.0232
0.00805
0.00535
0.00444
0.00124
0.0108
0.00436
0.00167
0.00385
<1.27>
Constants for this Data Set: b1 = 1.004,
SQ = 0.100 mg/A, AS =0 .090 mg/fc, Sin (S^S) =2.303
-------
• 0.000398
a(on TOG) - (2'33^j'14) = 0.0425
E - Zt - (0.375) (171)
bv ~ 60" ~ 60 - ~ 1<
-
AS (10 ) (0.090 x 10 J)
Ev 1.069
-
A (TOC) 10 (4.14 x 10
/ A TOC = 60(ATOC) = -
|TOC B E t ' B^, I t^ (227) (0.375) (171
= / A TOC = 60(ATOC) = - (6Q) (4-14> - T = 0.98 xlO~4
r| '
Values of these parameters for all conditions of chloroform in
lake water are summarized in Table 19.
5. Bromodichloromethane Destruction in Lake Water (A)
Example;
i
I = 0.375, (mB/V) = 1.33, t = 62.2 min,
S = 0.100 mg/A, AS = 0.090; TOCQ =4.3,
(TOC/TOC ) = 0.305, ATOC = 0.695(4.3) =2.99
B = (1.33) (62.2) = 82.7
C-
b (on S) = (B /AS) = (82.7/0.090) = 920
3. C
ba (on TOC) = (B^ATOC) = (82.7/2.99) =27.7
£n So/S 2 303
K (on S) = - °— = ,. *:,Q~ T/-.X = 0.0743
uv ( B/B°) (0.375) (82.7/1)
135
-------
TABLE 19. SUMMARY DATA ANALYSIS, CHLOROFORM IN LAKE WATER
(Tl
- Based on S -
•
I
0.189
0.375
""B
v
1.33
0.775
0.454
0.260
0.113
1.33
0.775
0.488
0.260
0.110
t
(min)
182
235
298
376
510
171
216
261
329
431
B
c
242.1
182.1
135.3
97.8
57.6
227.7
167.7
127.3
85.4
47.4
,TOC>
TOC,-,
0.17
0.20
0.26
0.35
0.52
0.037
0.053
0.084
0.160
0.340
b
a
2,680
2,024
1,504
1,086
640
2,522
1,864
1,414
948
526
K
u
0.0503
0.0669
0.0900
0.125
0.257
0.0270
0.0366
0.0482
0.0719
0.130
a
(X 105)
7.46
9.94
13.4
18.5
31.4
39.8
10.8
14.2
21.2
38.2
E
V
(XIO^J
3.19
4.11
5.22
6.58
8.93
1.19
7.50
9.06
11.4
14.5
b
a
67.8
52.9
42.5
35.0
27.9
54.8
41.2
32.3
23.6
16.7
<0.0903>
- Based on TOC -
K
u
0.0387
0.0468
0.0527
0.0568
0.0601
0.0387
0.0467
0.0519
0.0572
0.0607
<0.0510>
a
0.0344
0.0440
0.0548
0.0667
0.0833
0.0424
0.0566
0.0721
0.0984
EVTOC
161
215
295
423
7
258
332
414
570
935
n
(x io4)
1.41
1.09
0.840
0.643
0.437
0.998
0.832
0.732
0.625
0.516
Constants for
this
Data Set
: b'
= 1.004
, s =o
.100 mgA
, AS =
0.090 n
ig/A, £n(S_
/S) = 2.
303
For TOC: TOCO = 4.3 mg/fc, b1 = 2.33
-------
£n 4.3/1.31
K (on TOC) = . = 0
u (0.375) (82.7/1) 0-
a= blAS= (0.7324) (0.090) . ^^
V—'
a(on TOC) = * 8^7 ' = 0.0842
It = (0.375) (62.2) _
v 60 60 ~ °'389
E
EvS = V = (0.389) (1000) = 4340
AS(10-3) °-090
= 0.389(1000) _
vTOC 2.99
ATOC (2.99)
nTOC = B E t = (82.7) (0.389)62.2 = °-00149
C V
Values of these parameters for all conditions are summarized in
Table 20.
6. Simultaneous CF, BCM, TCE and TOC Destruction
A composite run was made to determine the simultaneous dis-
appearance of CF, BCM, TCE and TOC in Lake Water (A) at a
specified set of run conditions. The results are presented in
Table 21. it will be noted that for 90% removal of CF and BDty
-99% of the other TOC was removed. It appears that the UV-
intensity was beyond the optimum, permitting greater reduction of
other TOC than is necessary, cf. lower I-value, to get the
desired 90% reduction of the micropollutants.
137
-------
TABLE 20. SUMMARY DATA ANALYSIS, BROMODICHLOROMETHANE
DESTRUCTION IN LAKE WATER (A)
to
CO
•
T / \
J- I v )
mgO
watts/£ (—=—
mm
0.189 1.33
0.80
0.53
0.30
0.375 1.33
0.80
0.53
0.30
t
) &nin ) (
93.6
123.8
148.6
178.5
62.2
70.4
73.5
80.5
B
c
?*>3}
' £
124.5
99.1
78.7
53.6
82.8
46.3
40.0
24.2
,TOC %
o
- Based on S -
b K
a u
a
Ev
ba
- Based on TOC -
Ku
EVTOC
n
(X 103) (X 10^
0.407
0.440
0.471
0.572
0.305
0.379
0.466
0.607
1383
1101
874
596
920
626
444
269
0.0979
0.123
0.155
0.227
0.0742
0.109
0.153
0.254
<0.149>
0.530
0.665
0.838
1.230
0.796
1.171
1.648
2.724
3276
4333
5201
6248
4340
4889
5243
5590
48.8
41.1
34.7
29.1
27.7
21.1
17.4
14.3
0.0382
0.0438
0.0506
0.0551
0.0383
0.0460
0.0509
0.0550
<0.0472>
0.0497
0.0567
0.0672
0.0800
0.0842
0.1105
0.1340
0.1627
116
162
206
306
130
165
205
298
0.740
0.500
0.410
0.340
1.49
1.53
1.61
1.72
Constants for
this Data
Set:
bf = 0.7324, S
_ = 0.100
mg/Jl,
AS = 0
.090,
tn S /S -
2.303
For TOC: TOCQ = 4.3
, b1 = 2.33
-------
TABLE 21. COMPOSITE RUN, CF, BCM AND TCE IN LAKE WATER (A)
U)
vo
t
(min)
0
4.75
9.67
15.5
20.0
25.0
30.3
35.4
41.4
45.6
51.2
5$?
0
47.5
96.7
155
200
250
303
354
414
456
512
c/co KU (c/s)
1.00 - 0.0222
0.700 0.0189 0.0450
0.484 0.0189 0.0935
0.253 0.0223 0.181
0.178 0.0217 0.349
0.126 0.0209 0.756
0.0969 0.0194 1.90
0.0761 0.0183 4.70
0.0641 0.0167
0.0433 0.0173
0.0404 0.0158
<0.0190>
,TOC . A r
Imnr ' A expL-
J.ui_0 AO
(DA>0) KU (D/S) (E/EO) KU (E/s)
1.000 0.0222 1.000 0.0043
0.733 0.0165 0.0475 0.123 0.111 0.00158
0.518 0.0171 0.100 (0.0143) 0.111 0.00054
0.257 0.0221 0.184
0.166 0.0226 0.326
0.115 0.0218 0.930
0.0814 0.0209 1.60
0.0635 0.0196 3.92
-
-
- -
<0.0201> 0.111
OHT7QTT3 1 — QW> r_n mi^ma T
. UJ/o-Lii J exp L U • U-LDU.t>(2 J
\ft
(A/AQ) a
1.000
0.490 0.108
0.234 0.079
0.0977 0.058
0.0498 0.048
0.0235 0.039
0.0106 0.0330
0.0049 0.0282
Gas Flow = 0.666 A/tain @ 23QC, i^/V = 10 rag/A min, I = 0.397 watts/£
CQ(CF) = 0.100 mgA, DQ(BCM) = 0.100 rag/A, EJPCE) = 0.020 mg/Jl; A (TOG) = 4.30; S = 4.51
-------
7. THM Precursor Destruction in Lake Water (B)
Example :
I = 0.196 watts/A, m0/V = 2.09, t = 60 min,
B
THMPQ = MQ = 0.352 mg/A, M/MQ = 0.04; TOCQ =8.7 mg/fc,
b1 = 3.98, TOC/TOCo = 0.416
B = (2.09) (60) = 125
C*
b = B /AM = 125/0.96(0.352) = 370
3. C
b (on TOC) = B /ATOC = 125/0.584(8.7) =24.6 mgO_/mgATOC
3. C .3
£,n M /M
Ku(onH) - i ° - (0J^125) - 0.313
£n TOC /TOC
K (on TOC) = - - - = 0.0220
I Bc/B°
a (on TOC) = = (3.98)0 584(8.7) =
B 125
watt hrs/£
E M
V
AM(10 ) (0.96) (0.352) (10 )
0.196 = 38 6 (watt nrs'
(0.584) (8.7)10~3 g
ATOC (0.584)(8.7)
nTOC EcEvt (125) (0.196) (60) U
Values of these parameters for all conditions are summarized in
Table 22.
140
-------
TABLE 22. SUMMARY OF DATA ANALYSIS, THMP DESTRUCTION IN LAKE WATER (B)
I
"B
( V )
(watts/& ) (mgoy &
min
0.096
0.196
0.40
2.23
4.00
6.36
2.09
4.18
6.36
2.18
4.09
6.18
t
(min)
60
60
45
60
45
30
45
30
20
B
c
(M/MQ)
b
a
K
u
mgAM?
140
240
286
125
188
191
98.1
123
124
0.09
0.09
0.01
0.04
0.10
0.16
0.03
0.12
0.14
466
799
875
370
633
689
306
424
437
0.179
0.105
0.168
0.131
0.0625
0.0490
0.0894
0.0431
0.0396
<0.0963>
^M
(watt
hrs/g)
320
320
220
580
619
354
937
689
470
,TOC .
o
0.516
0.436
0.412
0.416
0.429
0.530
0.415
0.521
0.745
a
0.120
0.081
0.071
0.162
0.105
0.085
0.171
0.135
0.071
EVTOC
22.8
19.6
14.1
38.6
29.6
84.0
59.0
48.0
60.1
n
0.00522
0.00355
0.00552
0.00346
0.00399
0.00728
0.00384
0.00565
0.00671
Constants for this Data Set: THMP=M; = 0.330 mg/£
For TOC: TOCQ = 8.7 mgA; b' = 3.98
-------
C. REACTION RATES FOR MIXTURES
Examination of the data in Tables 16 to 22 incl. indi-
cates a substantial difference in the rate constants in lake
water compared to in purified water. Also, a substantial
reduction in other organic carbon is necessary to reduce the
micropollutants to the desired level of 90% removal. Obviously,
all oxidizable species present are oxidized to some extent and
thereby exert an ozone demand. By choice, the experimental work
tracked the disappearance of TOC, the micropollutants, and the
THM precursors, but not any inorganic species.
A model is needed which will account for the simultaneous
disappearance of TOC as well as the other species of primary
interest. Tables 23 and 24 present the laboratory data
obtained for TOC reduction. The most complete set is for Cross
Lake water giving values for ozone supply rate as well as at
three levels of UV-intensity. Several striking features of the
data will be noted.
1) - the 03 utilization values, a, are quite low, and are
not constant at different O3 supply rates. a-values
increase with increasing UV-intensity, but decrease
with increasing ozone supply rate. Concerning the
latter, it is true that increasing the O-j supply rate
without limit would lead to decreased a; however, over
the operating range used, it is not clear, at the out-
set, why this decrease should occur. Mass transfer
calculations presented later are revealing.
2) - the rate constants, calculated by £n(A /A)/(B /B°) vary
somewhat for different supply rates, primarily because
the ozone utilization is not the same.
Using the two sets of data, the constants were determined
for the following functions:
K
:) (1 +
u
K
(lOa)
(10b)
and for the reactor used, a * a (1 +
o
K
n
i)
(lOc)
142
-------
TABLE 23. TOG1DISAPPEARANCE, CROSS LAKE WATER
1
0.00498
0.00345
0.00242
<0.00372>
0.00526
0.00321
0.00269
<0.00372>
0.00817
0.00411
0.00325
<0.00518>
K+/K = 5.54;
u' o
= 0.0380, 2
a
A
^=•4- n A n\
lat — — — U.4U;
0.0385
0.0372
0.0254
<0.0377>
0.113
0.0783
0.0549
<0.0820>
0.119
0.0728
0.0610
<0.0843>
0.185
0.0932
0.0739
<0.1174>
n = 0.842,
\ = TOC
And the differential rate equation becomes
dt
A
,+
-A=
VB
VB
(10d>
143
-------
TABLE 24. TOC DISAPPEARANCE, LE.WISVILLE LAKE WATER
I B
C
0.084 20
40
52
0.189 20
60
100
140
0.375 20
60
100
120
(A/AQ)
0.915
0.835
0.800
0.800
0.625
0.340
0.255
0.620
0.257
0.127
0.092
<0>
0.0425
0.0413
0.0385
<0.0408>
0.100
0.0626
0.0661
0.0533
<0.0705>
0.240
0.124
0.0875
0.0758
<0.132>
Jin A /A
B /B
0.00444
0.00450
0.00427
<0.00440>
0.0112
0.00783
0.0108
0.00975
<0.00990>
0.0239
0.0226
0.0207
0.0199
<0.0218>
£n A /A
^aB /B° ^
0.105
0.109
0.111
<0.108>
0.112
0.125
0.163
0.183
<0.146>
0.0996
0.183
0.236
0.262
<0.195>
Constants: K = 0.0010; KU = 47.7; n = 1.00; r
K~~
A ETOC u
0.997; ao =0.048
The effect of UV-input shows up in the model as the K /K
ra.tio and the exponent "n" on the intensity. One would expect?
K /K to be several-fold to many-fold greater than unity, and
11H" to be in the range 1/2 to 3/2.
A composite model for all organic components including TOC
can be constructed in the following manner; the components are
defined for convenience as,
144
-------
S = total organic substrate components
A E other organic carbon components
B E ozone
C E chloroform (CF)
D E bromodichloromethane (BDM)
E E tetrachloroethylene (TCE)
M E trihalomethane precursors (THMP)
Then for CF, BCM and TCE in Lewisville Lake water,
S =A+C+D+E (lla)
(4.52 mg/£) (4.3) (0.100) (0.100) (0.020)
Assuming simultaneous disappearance of all components,
S = A + C + D + E (lib)
-S - KA(mB/VB°)(A/A°) + KC(mB/VB°)(C/C°) + KQ(mB/VB°
KE(mB/VB°)(E/E°)
-S = (mB/VB°)[KA(A/A°) + KC(C/C°) + KD(D/D°) + KE(E/E°)]
Dividing by S,
S ' A7Bo. rKA A . KC C , KD D KE B-, (lid)
- -S = (VVB > C^ S + ^ S + ^ S + ^ S]
Integrating,
S A
O
Hn &) = (B/B°) [(-
3 c A°t -o
and as an approximation,
v +-
f^dt + ...] dig)
where the total rate constant KA breaks down into the two parts,
for ozone only and ozone-UV, as shown by Equation lOb.
145
-------
By the mean-value theorem, the mean value of C/S over the
time interval involved is,
O CO
Obviously, if the rate of disappearance of S is greater than C,
then the ratio C/S will increase, going from a low initial value
to a high value, perhaps approaching unity. Actual values are
shown in Table 21 for the three components.
The final model involves three equations; one for the total
disappearance of S, a second for the disappearance of the
component requiring the longest time to attain the desired
reduction, and the third is Equation lOb, giving:
C
£n(^°-) = KC(I) (Bc/A°B°) (11 j
In order to use the equations, values for the rate constant
ratios and concentration ratios are required. For the micro-
pollutants values are presented in Table 25. It will be
noted that the rate constant ratios are for the most part
substantially greater than unity; however, since the concen-
trations are very low, the rate of removal is low cf. TOC
removal. The rate limiting component is chloroform, requiring
60-90% removal of TOC.
For removal of THMP (EM) the equation is,
A 00 K S
and the values of the rate constant ratio and the concentration
ratio are given in Table 26. In this case, it will be noted
that these ratios are essentially constant over a wide range of
ozone supply rates and UV- intensities. In order to get 90%
removal of THMP, approximately 50% removal of the TOC was
necessary.
146
-------
TABLE 25. REACTION PARAMETERS, MICROPOLLUTANTS IN LAKE WATER (A)
Component I B
c
TCE (E) 0.084 8.0 — > 27.0
0.189 5.2 — >19.4
0.357 3.6 — >14.8
0.397 97
CF (C) 0.189 58 — > 242
0.375 47 — > 228
0.397 250
BCM(D) 0.189 54 — > 125
0.375 24 — > 83
0.397 250
K <
KE °
0
0
0
KC o
0
0
KD °
0
0
.000
.221
.296
.0441
.0223
.0235
.00754
.0285
.0555
.00798
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
Comp
S
00257
00259
00260
00301
00309
00410
241
0138
0140
268
TOC
TOCo
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
92
87
84
30
30
13
004
47
44
004
0.
0.
0.
0.
0.
0.
0.
0.
0.
0.
KA
00411
00956
0195
0224
00947
0191
0224
00886
0179
0224
KA
27.
23.
15.
1.
2.
1.
0.
3.
3.
0.
0
1
2
97
35
23
337
22
10
356
0.0395
0.0906
0.128
0.130
0.0566
0.0674
0.0400
0.0634
0.123
0.0400
-------
TABLE 26. REACTION PARAMETERS, THMP IN LAKE WATER (B)
00
I (m
0.096 2
4
6
0.196 2
4
6
0.400 2
4
6
B/V>
.23
.00
.36
.09
.18
.36
.18
.09
.18
(M/M )
o
0.09
0.09
0.01
0.04
0.10
0.16
0.03
0.12
0.14
<0.087>
0
0
0
0
0
0
0
0
0
KM
.0172
.0101
.0161
.0257
.0123
.00960
.358
.0172
.0158
A/AQ
0.516
0.436
0.412
0.416
0.429
0.530
0.415
0.521
0.745
<0.491>
0
0
0
0
0
0
0
0
0
KA
.00498
.00345
.00242
.00526
.00321
.00269
.00817
.00411
.00325
VKA
3.45
2.93
6.65
4.89
3.83
3.57
4.38
4.18
4.86
<4.30>
0
0
0
0
0
0
0
0
0
<0
.0223
.0229
.0194
.0208
.0234
.0247
.0203
.0233
.0225
.0222>
a
0.120
0.081
0.071
0.162
0.105
0.085
0.171
0.135
0.071
THMP EM, = 0.330 mq/l, TOG = A; A = 8.7 mg/&; M /A = 0.0379
-------
D. REACTOR DYNAMICS — MASS AND ENERGY TRANSFER
There are numerous aspects of "reactor dynamics" involved in
the scale-up from laboratory scale experiments to full-scale
design. In this instance, the laboratory work was carried out in
a 3.0 liter (0.106 ft3) liquid volume reactor; whereas the full-
scale reactors will be in the range of 100 to 1000 ft3 liquid
volume. For various gas and liquid flow rates, and stirring
rates, the following considerations enter into the dynamical
problems involved:
1) the mixing patterns, degree of mixing, and power input;
2) the mass transfer of O-. from gas phase to liquid phase;
3) the mass transfer of volatile components from liquid
phase to gas phase; and
4) the energy (UV-range) transfer across semi-transparent
boundaries into the liquid.
The reactor used in the experiments is described in the
first part of the report and can be compared to the standard
reactor configuration, described previously and used by others
[2,8]. A comparison of the dimensional ratios, laboratory
reactor vs. standard reactor is made in Table 27.
TABLE 27. COMPARISON OF REACTOR DIMENSIONS
Dimensional Ratio
H /DT
VDT
VDT
h/DI
b/Dj
HI/DI
HS/H
Average gas
path length
Lab
Reactor
1.
0.
0.
0.
0.
1.
0.
12
11
547
135
222
395
21
39
cm
Standard
Reactor
1.
0.
0.
0.
0.
1.
30
00
333
100
200
25
00
0
cm*
For engineering laboratory reactor [2,8]
149
-------
In the standard reactor, the gas is introduced through a
sparger located at the bottom of the reactor. Whereas, in the
lab reactor used, the sparger was located at approximately
0.50 DT above the bottom, thereby reducing the gas path length
significantly.
Also, the UV radiators were externally located for the lab
reactor used. Whereas, in a standard reactor of engineering
scale and above, the radiators are inside the reactor, mounted
axially in the horizontal plane quadrants, thereby giving
significantly greater UV-radiation transfer into the liquid.
However, for the lab reactor the photon transfer was actually
determined by directly measured actinometry experiments.
Mixing Variables
In order to produce the necessary mixing patterns for
complete mixing, good UV-energy transfer, and high interfacial
area for mass transfer, the following are required for a
standard reactor,
DlNp 4 7
ND(Reynolds number) = ( — ) = 10 -> 10 (12a)
K y
N (gas flow number) = (100 Q/D3N) = 1 •> 10 (12b)
p (Power input)= 1 -*• 10 horsepower/1000 gals (12c)
The Bates and Moo-Young [9,10] correlations can be combined to
obtain the power input for the above range of NR and N .
0.679 PTD^N3 n.
P = ( „ v —)N"'U^ f(N ), (HP/1000 gal) (12d)
f (N ) = 0.45 exp("°-413) + exp(-0.525 N ) (12e)
Example Calculation — values of the above parameters can be
calculated for a typical run of the lab reactor.
= °'266 ft' N=30°
V = 3.00/28.32 = 0.106 ft3.
150
-------
, (0.266)2(300/60)(62.2)
(6.38 x 10 )
N = 100(1.90 x IP"4) = Q>2Q2
g (0.266) J(300/60)
f(N ) = 0.45 exp() + exp.(-0. 525) (0.202) = 0.958
Cj U , £ U Z ' """' "
_ ,-(0.679) (62.2) (0.266)5(300/60)3-1 n ., , 4. 0.045, 0
P - L (32.2) (0.106) J (3<45 X 10 } I0
= 3.15 (HP/1000 gals)
The calculation indicates a low gas flow number, and hence gas
velocity, as compared to an engineering scale or full scale
reactor, which would operate at 10 to 100 times greater Ng-valu5S
Ozone Mass Transfer
Ozone, like oxygen, is relatively insoluble in- water.
Maximum liquid phase concentrations at room temperature depend
on the gas phase concentration and lie in the range of 1 + 20
mg/X,. Henry's law provides the equilibrium relationship,
XB - P
4
H , the Henry's law constant being of the order of 10
Therefore, the overall mass transfer coefficient,
!_ + -1— = i_ ; K = k d3b)
KL
is liquid side controlled; hence the ozone mass transfer can be
described by,
151
-------
For a given situation of flow, etc., equation (13c) determines
the upper limit of 0., transfer into the liquid. If the actual
supply rate, (mB/V) , is less, all will be transferred; but if
greater, some of the 03 will remain in the bubbles and bypass
the liquid, flowing out the top of the reactor.
Ozone mass transfer coefficients have been determined and
correlated by Yocum [11] for a standard reactor configuration of
engineering laboratory scale. His data gave the relationship,
. r oc 0.67 0.454 , Ib mols >
k a = 6.85 u P , ( — 5 - )
L g ft min,AX
which can be normalized by introducing u = 0.60 ft/min, and
P = 10 HP/1000 gals, giving, g
This equation is consistent with the correlation function found
by Prengle and students [12] for O~ mass transfer in an engi-
neering laboratory scale standard reactor.
(13f)
The mass transfer coefficients obtained for the lab reactor
in this project at various conditions are summarized in Table
28.
Example Calculation — assume the following set of conditions
for the lab reactor:
• •
G = 0.323 s£/min; w%O_ = 2.5, m_ = 3.99 mg 0,,/min
•j B J
(mB/V) = 1.33 (mg03/min£) . T = 23°C; H = 4.20 x 103 atm.
Other conditions are the san\e as previous example (p. 152) •
r* = (W% °3^ fMcg^ P • 106 = (0.025) (28.9) (1.0 x 106)
- -
100 Mw H - (18)(4-20
mgO
152
-------
9-56 x 10 3 _6 g mols 0.
= b 3
Y - - _6
-- (48) (55 5) = 3'59 x 10 b (
io D.:D v
. - _
mol water
n - 323 mA/min n „
g T^ - 2~ = 1'88 cm/min - 0.0617 ft/min
172 cm
Using Yocum's correlation,
ft minAX
=28.6 („ ? m°ls)
v£min AX '
We note from Table 2 8 that the value for the lab reactor
be approximately 5 -»- 6 .
Now we can calculate the 0., transferred and compared with
the supply rate,
(mB/V)L = (48 x 103)kLa(X* - X) = (48 x 103)(6)(3.59 x 106)
mgO_ mgO-, supplied
= 1.03 (-7—V) cf. 1.33( —. 5- )
mm A — mm A
k a X
Defining, ND (0_ mass transfer number) = ( .— \ (13g)
B 3 mB/V ;
Then, Nfi = (1.03/1.33) = 0.774
Since NB < 1, not all the ozone was transferred into the liquid
phase in the lab reactor, for this case.
Table 2 8 also shows the N -values for the cases listed
plus four calculated cases which fall in the range of the
operating conditions used, indicating that in numerous
instances mass transfer was not complete.
Ozone Utilization
Since we are dealing with a heterogeneous process, gas to
liquid transfer followed by reaction in the liquid phase,
reactor dynamics enter into the efficiency of utilization of
ozone. This fact has been amply demonstrated previously [13].
153
-------
TABLE 28. SUMMARY OF
MASS TRANSFER COEFFICIENTS AND TRANSFER NUMBERS
m
Gas
Rate
(m£/min)
25.0
200
200
200
200
1000
1000
°3
Rate
(mg/min Si )
•0 . 4.7
3
2
1
0
7
7
.82
.86
.75
.954
.2
.2
Some other calculated
323
300
300
666
1
6
6
10
.33
.00
.00
.0
u
g
( cm/min )
0.145
1.
1.
1.
1.
5.
5.
16
16
16
16
8
8
N a L
(rpm)
300 1.98
500 2.22
500 2.60
500 2.43
500 2.15
500
500
/g
0 i
4
7
5
9
8
5
28
L
mols. ,lb mols .
•nirfAX ft^minAX
.44 0.277
.75 0.483
.73 0.357
.95 0.621
.11 0.506
.30 0.331
.1 1.75
kL
1 f
0
1
0
1
1
0
4
*
aX
mgO,
)
min £'
.765
.34
.987
.71
.40
.913
.84
N
1.
0.
0.
0.
1.
0.
0.
B
76
350
345
977
47
127
672
cases -
1.
1.
1.
3.
88
74
74
86
300
300
300
300
(6
(12
(30
(30
*
.0)
*
.0)
*
.0)
*
.0)
1
2
5
5
.03
.07
.17
.17
0.
0.
0.
0.
774
345
862
517
Estimated from the above data.
-------
By decreasing mixing, for both 03 only and 03/UV reactions, one
can lower the reaction rate, and hence the 0^ utilization value,
a, for a given ozone supply rate.
Consider as a point of departure an 03-only experiment.
For a very well mixed complete mass transfer situation, defined
by,
N^ = 104 + 107, Ng = 2 -> 10, p = 5 + 10, NB > 1
"there will be an intrinsic a-value, determined only by chemical
reaction rate. In addition, as observed from the data, for UV-
input the value of a increases significantly. Consequently,
can be represented functionally as,
+
a = f(aQ, kLa, mB/V, I, dynamic variables) (13h)
By equations 13e, 12d and 13g, N contains all the dynamical
effects and can be expressed as,
a = a* N?(l + alm) (131)
o B +
K
where, n - 1/2 ->• 1 , m = 1/2 + 3/2, and a - (^)
J\
o
Volatile Components Mass Transfer
Under operating conditions of full-scale equipment, volatile
components will be partially desorbed or stripped from the
liquid. These components, e.g. chloroform, are much more
soluble than substances like oxygen and ozone; and therefore
have much smaller Henry's law constants. Consequently, the
overall mass transfer coefficient,
^ = H_ + I- = * - K =k (14a)
Kg kL kg kg g g
is essentially gas side controlled. The mass transfer can be
described by,
p*
s = - = 2 (14b)
Data were obtained for chloroform using the lab reactor at
different gas velocities and stirring rates. The results are
155
-------
summarized in Table 2 9. lines 1-6 incl. There is a fair
amount of scatter in the mass transfer coefficients, but they
can be correlated by functions similar to equations (13e) and
(13f ) ; giving,
i a on in~6/ ^ \1-05 , P \° • 50 /g*GF , ,, , .
kga = 6.20 x 10 (^-|Q-) (^) , () (14c)
The magnitude of the values appear to be low, line 7 in Table
29 is a run obtained by Houston Research, Inc. [13] in an
engineering laboratory scale reactor which gives a coefficient
a.10.6 times greater.
Therefore, for large scale reactors, the following
function is more realistic:
^ mm ft
Example Calculation; Estimate the loss of volatile components
by stripping by gas flow through a continuous flow 2000 ft
stirred tank reactor, operating at 26°C and with a residence
time of ten minutes. Initial concentrations are: CF and BCM =
0.100 mg/A each and TCE = 0.020 mg/£ . The total removal of the
compounds of interest by oxidation and stripping is 90%.
TABLE 29. CHLOROFORM MASS TRANSFER DATA
Gas Rate ug
N
( Vniin) (cm/min) (ft/min) (rpm)
1) 0.200 1.16 0.0381
2)
3) 0.660 3.84 0.126
4)
5) 1.00 5.81 0.191
6)
300
500
300
500
300
500
3
5
3
5
3
5
NR
.45xD
.75x30
.45x30
.75x30
.45x10
.75x10
N
g
4
4
4
4
4
4
0.125
0.
a
0
075
413
.248
0.625
0
.375
p
(HP/1000 (—. — p)
gal) min
3.
15
3.
15
3.
15
14
.6
20
.0
14
.2
0
0
0
0
1
2
.195xl06
.499xl06
.733xl06
.781xl06
.82x!0"6
.02xlO~6
7) 4.00 9.36 0.307 700 7.40xlQ4 1.22 8.94 318x10 6
156
-------
ft/min and the
k a = (4.098 x 10 ) ( '3 ) (-^5.^ - a 07 v i n~6 / lbs N
cr ^ .r>. j-v '*fifiiVMn' — o.o/xj.u (—: ^-i
y U.bU 10 u vmin ftj'
^
P Q
0
= (8.87x10"
• • -v-h ()• ' % i nn '
lbs. removed
UV-Energy Transfer
The chemical reactions involved are driven by the UV energy
input to the liquid. An energy transfer process is involved and
must be dealt with, both for the lab reactor used and scale up
to full size.
UV-energy (180 to 400 nm) is transferred into the liquid
phase from the UV-radiators located in the reactor, or as in the
case of the lab reactor, located external to the reactor.
Commercially available UV radiators have a total wattage rating,
which measures the power requirement, but only a fraction (15-
25%) of the total power (w ) is converted to UV power (w ) , the
remaining being primarily infrared and dissipated as hea^.
The UV-energy must pass through radiator wall, quartz sheath,
and reactor walls, whatever the case may be, and thence into the
liquid phase — the rays being of different path lengths (L).
The integrated UV intensity in the liquid can be represented by
a transfer equation of the form,
w ' nw *
I = (_U) . (_J!) . K (15a)
W V V
(watt UV in , " „., f , watts UV in liquid ,
liquid/*) lT£Ht™s/ir Wt UV from radiators'
w nw ,T
_ / U) / D.) (13) (i - e ) (15b)
vw V aL
n
Kv being the UV-energy transfer coefficient accounting for geom-
etry (g, L) and attenuation factors (T,a). The geometric factor
g, is determined by the location of the radiators; if located
inside the reactor g = 1, but if located externally,
157
-------
g(r,R) = (l - sin"-^) (15c)
The attenuation coefficient for quartz wall [14], T = 0.90, and
for aqueous media, evaluated from experimental data [15],
a = 0.17 to 0.21.
Consequently for a given reactor, UV-radiators, and
geometry, the UV intensity in the liquid can be calculated
using equations 15b and 15c. A calculation for the lab reactor
is illustrated by the following example.
Example; Calculate the UV-intensity in the liquid for the lab
reactor using 4-14 watt bulbs mounted externally, and compare
with 1-bulb immersed in the liquid. UV watts equal 18.8% of
nominal watts, = 10.5 cm; assume a = 0.175, T = 0.90.
2
g =
K* - !2n-e-aL) - o-90(Q-260) ri-e"°-175(10-5)i - o 107
Kv ~ aT(1 e } ~ 0.175(10.5) Li e J ~ 0'107
I = K (^i) (-^) = )0.107) (0.188) (4) (14)/3.0 = 0.375 watts/Jl
n
For 1-bulb (rather than four) immersed in the liquid.
o = 1- K* = (1)(1) ,-0.175(7.4) _ 0
g lf v (0.175) (7.4) ! e ~ °
I = (0.561) (0.188) (1) (14J/3.0 = 0.492 watts/£, approximate^
5-1/4 times greater.
Therefore- UV-transfer coefficients have the following
approximate values for different scale reactors:
Reactor Scale
Lab
Engr . Lab
Industrial
Radiators
External
Internal
Internal
*
Kv
0.10
0.50 - 0
0.40 - 0
.60
\
.60
158
-------
E. CONCLUSIONS FROM THE DATA ANALYSIS
As a result of a detailed engineering analysis of the
laboratory data, the following conclusions can be drawn.
1. Destruction of tetrachloroethylene in purified water showed
relatively high rate constants, but lower O3-utilization
values than would be expected, even with highest UV-intensity.
On the other hand, the rate constants in lake water were
5-10 fold less than in purified water, but the do-utiliza-
tion was somewhat greater when other TOG destruction was
included; however, the absolute magnitude of total Oo-
utilization was still low.
2. Destruction of chloroform in purified water showed factor of
10 lower rate constants than for TCE, and 03-utilization was
very low. In lake water, the rate contants were further
reduced by another factor of 10, but total 03-utilization,
including total TOC removal, was higher. However, the
absolute magnitude of (^-utilization was quite low. Sub-
stantial removal of TOC, 50-90%, occurred in order to get
90% removal of CF.
3. Destruction of bromodichloromethane in lake water showed low
rate constants, low O-j-utilization values, and substantial
removal of TOC, 40 to 70%, to obtain 90% removal. 03-
utilization was improved when the TOC destruction was
included, but on an absolute basis the values were still
quite low.
4. Destruction of trihalomethane precursors in lake water
showed low rate constants, approximately the same level as
for chloroform, and low 03~utilization values including TOC
removal. TOC removal amounted to approximately 50% to
get 90% removal of THMP.
5. The effect of UV-input is substantial in producing higher
rates and greater 03~utilization for both micropollutants
and THMP.
6. Insofar as design of full-scale equipment is concerned, the
reaction rate data for the micropollutants studied in
purified water provide little information. The rate con-
stants are many-fold less in lake water, which is to be
expected since numerous components are competing for
oxidizer, and furthermore the components, other than the
micropollutants, are in greater concentration. Therefore,
the crucial factor is the extent to which other TOC must be
removed in order to get 90% removal of the micropollutants.
159
-------
7. The scale of the laboratory reactor was too small to provide
data which can be used to scale up to commercial size for the
aspects of: O-^-mass transfer, volatile components mass
transfer, and UV-energy transfer. In particular, the
location of the sparger at 0.50 Dm, and the actual small gas
path length, 12 cm, did not provide sufficient contact time
for transfer.
8. In general, the dynamical mixing characteristics of the lab
reactor appear to be approximately satisfactory. NR and P
fall in the right range, but N -values are low by a factor
of 5-10; this coupled with thegposition of the sparger,
adversely affect the mass transfer coefficients.
9. The mass transfer coefficients in the lab reactor were low
by approximately a factor of 5-10, and in fact the calcula-
tions indicate that for many runs not all of the 0.,-supplied
was transferred into the liquid, accounting in part for the
low values of (^-utilization.
10. The mass transfer coefficients for volatile components, e.g.
chloroform, appear to be low by at least a factor of 10 in
the lab reactor.
11. The UV-energy transfer, coming from radiators on the outside
of the lab reactor, was substantially less than would be
achieved by radiators located inside the reactor, as in a
full-scale reactor.
F. SIZING METHODS AND CASES
For purposes of sizing and costing a range of full-scale
units, the following six (6) cases were defined:
Cases A-l, A-2, A-3
(a) Matrix: Lake Lewisville, Texas
(b) Influent Micropollutant Levels:
CHC13, Chloroform: C = 100 yg/£
CHCl2Br, bromodichloromethane: D = 100
C2C14, tetrachloroethylene: E = 20
(c) % Removal for all Micropollutants: 90%
(d) Plant Sizes: Case A-l: 1 MGD
Case A-2: 10 MGD
Case A-3: 50 MGD
160
-------
Cases B-l, B-2, B-3
(a) Matrix: Cross Lake, Louisiana
(b) Influent THM Precursor Level: Ambient
(c) % Removal of THMFP: 90%
(d) Plant Sizes: Case A-l: l MGD
Case B-2: 10 MGD
Case B-3: 50 MGD
For a given case defined by:
water flow rate, influent concentrations of TOC and
micropollutants or THM precursors, and required removal,
a procedure was used involving a sequence of steps to determine:
the ozone required, and hence the size of the ozone generation
unit, the reactor system size, and the energy required to
operate the system. In order to achieve the advantages of
higher average concentration of oxidizable components, a multi-
staged reactor system of four completely mixed stages in series
was used. Also, the combination of the four stages plus choice
of the dynamical variables to give complete mass transfer of
the 03 supplied, were incorporated into the design in order to
achieve high ozone utilization, a > 0.50.
The following mathematical relations — a total of some 24
equations -- were solved simultaneously, by numerical iteration
until the desired convergence was attained.
1. mass balances of the oxidizable components, ozone,
and air;
2. reaction rate equations for the disappearance of the
TOC and the individual compounds of interest;
3. mass transfer equations for ozone and volatile compounds;
4. UV transfer equations from radiators into liquid phase;
and
5. the mixing dynamic equations.
The iteration loop is started by making a first approxi-
mation of the design, followed by a calculational check of all
the conditions that must be satisfied. Since there are
numerous engineering solutions to the design problem, because
of trade-offs, the final optimization is based on an "operating
cost" (including amortization of the capital cost) object func-
tion. Time and funding limitations do not permit presentation
of the generalized procedure for the optimization.
161
-------
Table 30 presents the sizing results for Series A, and
Table 31 the results for Series B.
G- TREATING UNIT CONFIGURATION
The process for both cases, A and B, is as shown in Figure
38.
The influent water is pumped to Reactor Rla and by gravity
flow goes to Reactor Rib from whence, after treatment, it is
discharged for utilization.
Ozone is generated from either air or 02 and is fed to the
bottom of each reactor for best ozone utilization.
The off-gas from the reactors is fed to a gas-liquid
separator and in the case of oxygen-generated ozone, the excess
oxygen is recycled to the ozone generator.
A transfer pump is provided to feed the reactors from the
storage reservoir.
A flow indicator regulator is provided on the input to the
reactors to regulate the influent flow.
Ozone monitors are provided at various points in the system
to insure proper ozone utilization.
Each of the series reactors, as shown in Figure 38, is
continually stirred and contains the UV lamps, fully immersed
in quartz wells to insure proper UV energy transfer.
The unit, as pictured is easily adaptable to skid-mounting,
thereby reducing the on-site construction costs for small
communities.
H. CAPITAL AND OPERATING COSTS
Once the ozone requirements had been determined, the process
configuration developed and the reactors sized, the following
bases were used to estimate the capital and the operating costs.
PCI-Ozone Corp. was contacted and provided capital
estimates for the ozone generation equipment in the various size
ranges as well .as air preparation equipment or oxygen recycle
equipment, depending upon the configuration.
Estimates for reactors with mixers were obtained from local
suppliers of comparable equipment to the petrochemical industry.
This source was selected since most manufacturers of water and.
wastewater treatment vessels do not provide sufficient mixing
horsepower for this application.
162
-------
TABLE 30. SUMMARY OF DESIGN CASES SIZINGS, SERIES A
en
u>
Parameter
1) Flow, MGD
2) 03req'd, (Ibs/day)
3) TOC + Micropol-
lutants
-Influent, (Ibs/day)
-Removal, (Ibs/day)
4) Air or O~ req'd,
(Ibs/day)
5) Reactors x Stages
6) Reactor Size, (each)
D' x V(ft3) x H'
7) Mixing req'd,
HP
KWH/day
Motor Size, HP
8) UV req'd,
KW
KWH/day
9) Pumping req'd,
HP
KWH/day
Motor Size, HP
A-l
1.00
113
(from air)
37.7
26.4
4492
2x2
5'k236xl2'
(+ 3 'skirt)
25.7
460
2(18)
22.3
535
12.5
224
15
A-2a
10.0
1193
(from air)
377
264
47,720
2x2
10'xl855
x24' (+ 3'
skirt)
200
3581
2(180)
178
4271
125
2238
150
A-2b
10.0
1193
(from 02)
377
264
2386
(make up)
2x2
=
=
=
=
=
=
=
=
=
A-3a*
50.0
5965
(from air)
1885
1320
239,000
5(2 x 2)
5 units
1000
17,905
10 (180)
1000
21,355
625
11,190
5(150)
A-3b*
50.0
5965
(from 02)
1885
1320
11,930
(make up)
5(2 x 2)
5 units
=
=
=
=
—
=
—
*5-10MGD units for this case.
Design "Basis: CFQ and BCMQ =
0.100 mg/£, TCEQ = 0.020 mg/£ , TOC =4.3 mg/£
90% removal of micropollutants, 60-70% removal of TOC.
-------
TABLE 31. SUMMARY OF DESIGN CASE SIZINGS, B-SERIES
CTi
Parameter
1) Flow, MGD
2) 0Q req'd (Ibs/day)
J
3) TOC + THMP
-Influent, (Ibs/day)
-Removal, (Ibs/day)
4) Air or 0_ req'd,
(Ibs/day)
5) Reactors x Stages
6) Reactor Size, (each)
D' x Vft3' x H1
7) Mixing req'd.
-HP
-KWH/day
-Motor Size, HP
8) UV req'd,
-KW
-KWH/day
9) Pumping req'd,
-HP
-KWH/day
-Motor Size, HP
B-l
1.00
255
(from air)
72.5
38.4
10,200
2x2
5*236x12 '
(+3 'skirt)
25.0
447
2(18)
22.3
535
12.5
224
15
B-2a
10.0
2550
(from air)
725
384
102,000
2x2
10'xl855x
24' (+ 3'
skirt)
200
3581
2(150)
178
4271
125
2238
150
B-2b
10.0
2550
(from 02)
725
384
5100
(make up)
2x2
=
=
=
~
=
-
=
=
=
B-3a*
50.0
12,750
(from air)
3625
1920
510,000
5(2 x 2)
5 units
1000
17,905
10(150)
890
21,355
625
11,190
5(150)
*
B-3b
50.0
12,750
(from 02)
3625
1920
25,500
(make up)
5(2 x 2)
5 units
=
=
=
=
=
•=.
=
~
*5-10MGD units for this case.
Design Basis: THMPQ = 0.330 mg/£; 90% removal of THMP; TOCO = 8.7 mg/Jl;
50-60% removal of TOC.
-------
ui
Water
0. or Air
Vent
0_ Monitor
.S)-*-
0,/UV Reactor
R-lo
0 /UV Reactor
o
R-2a
I t
c
' I)
i r
ff
/-\/7i
cm
. i
,_,
HI ii i
111!
i=v
i
i
i
j 0 Monitor
-r-r--"-="-m Other Locations
Effluent for potable use
[_j
. RM
Figure 38. Oxyphotolysis Process Unit for Drinking Water.
-------
UV equipment estimates were provided by Canrad-Hanovia, a
supplier of industrial UV lamps.
The estimates for the pumps were provided through Flow -
Quip, Inc. using Allis-Chalmers pumps.
The. estimates for ozone monitors were provided by PCI and
the instrumentation estimate was provided by local suppliers.
Piping, valves and skid-mounted installation, where
feasible, was estimated by local fabricators. Design, engi-
neering and fee estimates were based on local industry percen-
tages. A 5% miscellaneous and contingency fee was also included
in the estimated capital cost.
The elements of the estimated direct operating cost include
operating labor, electrical energy cost, equipment maintenance
cost, UV lamp replacement cost and oxygen, where necessary.
Fifteen dollars per hour was used for the estimated cost of
operating labor. The daily operating time required ranged from
4 hours/day to 16 hours/day depending upon the size of the
treatment unit.
A cost of $0.02/KWH was used to estimate the electrical
energy cost. The elements composing the electrical energy load
include ozone generation, UV lamp operation, mixing, pumps and
instrumentation.
Maintenance is estimated at 1.5% of the capital cost per
annum while UV lamp replacements are estimated on the basis of
a 4000 hour life.
Oxygen is estimated at $70/ton delivered to the site in
those cases where oxygen is used to produce the ozone.
The total operating cost also includes amortization of the
capital expenditure. A 20-year amortization period has been
used at 7% interest.
Tables 32 through 35 summarize the capital and operating
costs for both the A and B series cases.
Larger cases, beyond 50 MGD, were not sized and costed,
since it is believed that process economics better than 2X50 MGD,
for example, can be achieved by a different design more
appropriate to very large volumetric flow rate units. Such
designs are in preliminary stages of development.
Recently it has been proposed (16) that for purposes of
estimating capital and operating costs, a correlation better
166
-------
TABLE 32. SUMMARY OF CAPITAL COSTS (A-SERIES)
en
Flow (MGD)
A-l
1.00
O3 req'd (Ibs/day) 113
(from air)
Removal (Ibs/day
TOC + Micro-
pollutants
Equipment
Ozone Generators
Air Prep or 02
Recycle
Reactors w/mixers
UV Equipment
Pumps
O., Monitors
Instrumentation
Sub-total
Piping^ Valves,
Instl.
Design, Engineering,
Fee
Misc. & Contingency
TOTAL
26.4
$ 81,000
21,000
11,000
8,920
3,000
6,000
10,000
$140,920
90,000
41,565
13,624
$286,109
A-2a
10.0
1193
(from air)
264
$480,000
60,000
75,000
80,000
20,000
6,000
10,000
$731,000
382,750
200,475
65,711
$1,379,936
Case
A-2b
10.0
1193
(from 02)
264
$200,000
60,000
75,000
80,000
20,000
6,000
10,000
$451,000
212,750
119,475
39,161
$822,386
*
A-3a
50.0
5965
(from air)
1320
$2,400,000
300,000
375,000
356,000
100,000
30,000
40,000
$3,601,000
1,500,250
918,225
300,974
$6,320,449
A-3b*
50.0
5965
(from 02)
1320
$ 900,000
250,000
375,000
356,000
100,000
30,000
40,000
$2,051,000
1,012,750
551,475
180,781
$3,796,006
*5-10 MGD units for this case.
-------
TABLE 33. SUMMARY OF OPERATING COSTS (A-SERIES) $/DAY
Operating Labor
@ $15/hr.
Electrical Energy
@ ($0.02/KWH)
0- Generation
UV Lamps
Mixing
Pumps & Inst.
Sub-total
en Maintenance
03 @ $1.5% of capital/
annum
UV Lamp Replacement
Oxygen @ $70/ton
DIRECT OPERATING COST
AMORTIZATION
(20 years @ 7%)
TOTAL OPERATING COST
($/day)
TREATMENT COST
($/1000 gals)
A-l
$ 60.00
$ 22.60
10.70
9.20
4.48
$ 46.98
11.76
5.85
-
$124.59
41.94
$166.53
$ 0.16
A-2a
$ 60.00
$238.60
85.42
71.62
44.76
$440.40
56.71
46.67
-
$603.78
202.26
$8.06.04
$ 0.081
Case
A-2b
$ 60.00
$109.76
85.42
71.62
44.76
$311.56
33.80
46.67
83.51
$535.54
120.54
$656.08
$ 0.066
A-3a
$ 240.00
$1,193.00
427.10
358.10
223.80
$2,202.00
259.74
233.36
-
$2,935.10
926.42
$3,861.52
$ 0.077
A-3b
$ 240.00
$ 548.78
427.10
358.10
223.80
$1,557.78
156.00
233.36
417.55
$2,604.69
556.40
$3,161.09
$ 0.063
-------
TABLE 34.
SUMMARY OF CAPITAL COSTS (B-SERIES)
V£>
Flow (MGD)
03 req'd (Ibs/day)
Removal (Ibs/day)
TOC + THMP
Equipment
Ozone Generators
Air Prep or O2
Recycle
Reactors w/mixers
UV Equipment
Pumps
O~ Monitors
Instrumentation
Sub-total
Piping, Valves, Instl.
Design, Engineering,
Fee
Misc .& Contingency
TOTAL
B-l
1.00
255
(from air)
38.4
$118,400
26,600
11,000
8,920
3,000
6,000
10,000
$183,920
108,850
52,695
17,273
$362,738
B-2a
10.0
2550
(from air)
384
$1,140,000
180,000
75,000
80,000
20,000
6,000
10,000
$1,511,000
906,600
435,168
142,638
$2,995,406
Case
B-2b
10.0
2550
(from O-)
384
$440,000
100,000
75,000
80,000
20,000
6,000
10,000
$731,000
438,600
210,528
69,006
$1,449,134
*
B-3a
50.0
12,750
(from air)
1920
$5,280,000
540,000
375,000
356,000
100,000
30,000
40,000
$6,721,000
3,965,390
1,923,550
630,497
$13,240,437
*
B-3b
50.0
12,750
(from 0«)
1920
$2,100,000
480,000
375,000
356,000
100,000
30,000
40,000
$3,481,000
2,053,790
996,262
326,553
$6,857,605
5-10 MGD units for this case.
-------
TABLE 35. SUMMARY OF OPERATING COSTS (B-SERIES) $/DAY
-j
o
Operating Labor
@ $15/hr
Electrical Energy
@ ($0.02/KWH)
03 Generation
UV Lamps
Mixing
Pumps & Inst.
Sub-Total
Maintenance
@1.5% of capital/
annum
UV Lamp Replacement
Oxygen @ $70/ton
DIRECT OPERATING
COST
AMORTIZATION
(20 years @ 7%)
TOTAL OPERATING COST
($/day)
TREATMENT COST
B-l
$ 60.00
$ 51.00
10.70
8.95
4.50
$ 75.15
14.91
5.85
-
$155.91
53.17
$209.08
$ 0.20
B-2a
$ 60.00
$510.00
85.42
71.62
44.75
$711.79
123.10
46.67
-
$941.56
439.05
$1,380.61
$ 0.138
Case
B-2b
$ 60.00
$234.60
85.42
71.62
44.75
$436.39
59.55
46.67
178.50
$781.11
212.41
$993.52
$ 0.099
B-3a
$ 360.00
$2,550.00
427.10
358.10
223.80
$3,559.00
544.13
233.36
-
$4,696.49
1,940.72
$6,637.21
$ 0.133
B-3b
$ 360.00
$1,173.00
427.10
358.10
223.80
$2,182.00
281.82
233.36
892.50
$3,949.68
1,005.16
$4,954.84
$ 0.099
($/1000 gals)
-------
than,
$/1000 gals treated vs. MGD
is S/lb oxidizables removed vs. Ibs removed/day.
The method is particularly applicable to chemical oxidation
since the oxidant required is directly related to the amount of
material to be oxidized per unit time.
Figure 39 presents correlation of the costs for Series A and
B on the suggested basis. By comparison, data for the
OXYPHOTOLYSIS1™ process applied to industrial wastewater gives
values of $0.75 to $1.00/lb removed, for cases in the range of
100 and greater Ibs removed/day. These lower treating costs
are to be expected, since in industrial cases, Oo-utilization
approaches 100%.
Sensitivity of the Cost Estimates
Examination of Tables 3 3 and 3 5 indicates that best
economics result when the ozone is generated from oxygen rather
than air for unit processing 10 MGD and greater; this fact might
also be true for units of 5.0 MGD. In order to determine the
sensitivity of the operating cost estimates, four additional
cases were estimated, designated A-2bf, A-3b', B-2b', and B-3b',
assuming that ozone utilization was only 35%. The results are
presented in Tables 3 6 and 37. In the case of the A-case
comparisons,the treating cost increased by approximately 15%,
whereas for B-case comparisons the increases were approximately
40%, but for all cases the estimates are less than $0.15/1000
gal. treatment cost.
I. ENGINEERING CONCLUSIONS AND RECOMMENDATIONS
In addition to the conclusions in Part E, and as a result
of the process engineering sizing and costing, the following
engineering conclusions and recommendations are offered:
1. The ozone/UV photo-oxidation process can be used
effectively to reduce micropollutants and trihalo-
methane precursors to the desired limits.
Simultaneously, TOC will be reduced and disinfection
will be accomplished.
2. The cost estimates indicate that daily operating costs,
including amortization of capital equipment and
interest on borrowed capital, are in the range of $0.06
to $0.15 per 1000 gals of water treated, for 10-50 MGD
units using ozone generated from oxygen.
171
-------
i
to ^
O -W-
o ~-
g
Q.
O
20,000
ropoo
5,000
3,000
(/>
o
z
Ul
Q.
O
ipoo
10.0
5.0
3.0
1.0
10
100
O O SERIES A
-I-A SERIES B
1000
10,000
Ibs. removed/day
Figure 39. Correlation of Capital and Operating Costs (Assuming 50-60% 03 Utilization).
-------
TABLE 36. OPERATING COST ESTIMATE SENSITIVITY (A-SERIES)
to
1)
2)
3)
4)
5)
6)
7)
8)
Flow (MGD)
03 Required (Ibs/day)
Removal (Ibs/day)
Capital Cost ($)
Direct Operating Cost
(S/day)
Amortization (20 years @ 7%)
Total Operating Cost
($/day)
Treatment Cost
(S/1000 gal)
($/lb removed)
A2b
10.0
1,193
264
822,386
535.5
120.5
656.0
0.066
2.48
A2b'
10.0
1,705
264
1,133,000
571.0
165.9
736.9
0.077
2.79
A3b
50.0
5,965
1,320
3,796,000
2605.7
556.4
3161.1
0.063
2.39
A3b'
50.0
8,524
1,320
5 ,426,000
2846.2
795.1
3641.3
0.073
2.76
-------
TABLE 37. OPERATING COST ESTIMATE SENSITIVITY (B-SERIES)
1)
2)
3)
4)
5)
6)
7)
8)
Flow (MGD)
O3 Required (Ibs/day)
Removal (Ibs/day)
TOC + THMP
Capital Cost ($)
Direct Operating Cost ($/day)
Amortization ($/day)
(20 years @ 7%)
Total Operating Cost ($/day)
Treatment Cost ($/1000 gal)
($/lb removed)
B2b
10.0
2,550
384
1, 449,134
781.1
212.4
993.5
0.099
2.59
B2b'
10.0
4,371
384
2,461,000
1057.7
360.7
1418.4
0.142
3.69
B3b
50.0
12,750
1,920
6,857,605
3949.7
1005.2
4954.8
0.099
2.58
B3b'
50.0
21,854
1,920
11,740,000
5157.2
1720.8
6878.0
0.138
3.58
-------
3. A treating cost of $0.06 to $0.15/1000 gals for
combined micropollutant, trihalomethane and TOG
removal, plus disinfection appears reasonable.
4. If higher ozone utilization values can be achieved,
the capital and operating cost could be reduced
significantly; the latter by as much as 30-40%.
5. For raw water streams containing TOG > 10 mg/L , the
possibility of coupling the process with another pre-
TOC removal process to minimize total treating cost
may be desirable. For high TOG streams, each
situation should be examined individually.
In view of the foregoing conclusions, it is recommended that:
The process should be field tested at the 1.0 MGD, or
greater, level for at least one year, using a four-stage
reactor system unit, as described in this report, to obtain
actual data on performance (operating efficiency, ozone
utilization, etc.) and treatment costs. The unit could be
completely skid-mounted for use at several locations, if
desired.
175
-------
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189
-------
APPENDIX A
DESTRUCTION CURVES FOR FIRST-ORDER SUBSTRATES
Shown in Figures A-l through A-16 are the disappearance
curves for substrates which are first-order in substrate.
Points are actual experimental data while the solid lines are
the results of the curve fits given in Appendix C. All curves
are normalized-to the initial concentration of substrate,
100 ug/L in all uses.
190
-------
1.0
0 10 30 50
time, seconds
20 40 60
time, seconds
8O 100
1.0
.8
V6
.2
•vG
20 40 60 80 100 120
time, seconds
1.0
.8
.6
20 40 60 80
time, seconds
100
50 100 150 200
time, seconds
25O 300
Figure A-l. Tetrachloroethylene in purified water. UV intensity = 0.375 W/L.
-------
1.33 ing/J-min.
ID
NJ
0.775 mg/J-min.
0 50 IOO
time, seconds
1.0
.8
.6
0.23 mg/(-min.
SO 100
time, seconds
1.0
.4
.2
50 100 ISO 200 250 300
time, seconds
1.0
.8
S, -6
A
.2
0.46 mg/JL'min.
©
0 50 100 ISO 200 250 30O
time, seconds
O.ll mg/ ?-min.
SO 100
ISO 200 2SO
time, seconds
300 350 4OO 4SO SCO 550 600
Figure A-2. Tetrachloroethylene in purified water. UV intensity = 0.139 W/L.
-------
vo
0.42 mg
1.0
0.8
S 0.6
0.4
0.2
0 100 200 300 400
time, seconds
1.0
So 6
.4
O ZOO 400 0
time, seconds
0.09 mg/ jt min
200 400 600
time, seconds
0.20 mg/j? min
O 200 400 60O 800 IOOO
time, seconds
0 200 4OO 600 800 IOOO 1200 1400 I6OO 1800 2000 2200
time, seconds
Figure A-3. Tetrachloroethylene in purified water. UV intensity = 0.091 W/L.
-------
1.0
.5
1.0
D= 0. 24 mg/ min- L
I
500
time, seconds
1000
1500
.5
D= 0.36 mg/ min-L
, I
500 1000
time, seconds
1500
D= 0.63 mg / min- L
1.0
.5
_ 0
D = 0. 48 mg / min - L
J_
500
time, seconds
1000
Figure A-4. Chloroform in purified water. UV intensity = 0.417 W/L.
-------
1.0
\
.5
D= .0.73 mg/min. - L
500
time, seconds
1000
D = 0.48 mg/min - L
500
time, seconds
500
time, seconds
1000
vo
Ui
1.0
.5
D= 0.46 mg/L
i , , i
500
time, seconds
10
1000
D = 0.32 mg/min- L
Q
i . i
500 1000
time, seconds
1500
Figure A-5. Chloroform in purified water. UV intensity = 0.184 W/L.
-------
1.0
0
.5
D=0.33 mg / min- L
i i i
i i i i i i i i
500 1000
time, seconds
1500
.5
D = 0. 44 mg/min - L
I i I I I I l I l I l l l I I l
500 1000
time, seconds
1500
a\
D = 0.55 mg /min - L
I I I I l I I I __ I L .1
500 1000
time, seconds
D = 0.72 mg/min - L
0 i i i i i i i i i i
500 1000
time, seconds
1500
Figure A-6. Chloroform in purified water. UV intensity = 0.084 W/L.
-------
•D = 1.33 mg/L min
10 20
time, min
D- 0.5 mg/ L> min
o
REPLICATE RUNS
10 20 30 40
time, min
20 30
time, min
1.0
0.8
0.6
'So
0.4
0.2
D= 0.3 mg/Li min
1.0
0.8
0.6
S/
/So
0.4
0.2
10 20 30 40 50 60
time, min
0.15 mg/Li min
O
O
O
10 20
30 40
time, min
O
0
0
50 60 70
Figure A-7. Bromodichloromethane in purified water. UV intensity = 0.375 W/L.
-------
1.0
0.8
0.6
0.4
0.2
f 1.33 mg/L min
- o
Q
10 20 30
time, min
40
0=0.5 mg/L. min
20 30 40
time, min
10
20 30 40
time, min
H
U7
00
1.0 t
0.8
0.6
X
0.4
0.2
D= 0.3 mg/L min
10
20 30
time, mm
40 50
1.0
0.8
0.6
0.4
0.2
D- 0.15 mg/L. min
10 20
30 40
time, min
50 60
Figure A-8. Bromodichloromethane in purified water. UV intensity = 0.189 W/L.
-------
I.Or
0.8 -
0.6
'So
0.4
0.2
0
D= 1.33 mg/L min
10 20
time, min
30 40
1.0
0.8
0.6
o
0.4
02
0
D = 0.775 mg/L min
10 20
time,min
Qi
0.5 mg/L min
30 40
10 20 30 40
time, min
D- 0.3 mg/L min
10 20 30
time, min
40
1.0
0.8
0.6
s/s.
0.4
0.2
D= 0.15 mg/L min
10 20
30 40
time, min
50 60 7O
Figure A-9. Bromodichloromethane in purified water. UV intensity = 0.084 W/L.
-------
to
o
o
1.0
0.8
Qfi
0.4
0.2
D = 1.33 mg/L -min
1.0
.75
200
400 600 800
time, seconds
1000 1200
D= 0.775mg/L -min
1.0
0.8
0.6
0.4
0.2
D » 0.26 mg/L - min
1.0
.75
.25
O
D= 0.417 mg/L -min
200 4OO 600 800
time, ucondt
1.0
0.8
1000
ZOO
400 600 800 1000
time, seconds
1200
Q2
O
D = 0.096 mg/L - min
O
200 400
600 800 1000 "1200
time, seconds
1400 1600
200 40O 600 800 1000 1200 1400 1600
time, seconds
Figure A-10. Tetrachloroethylene in Lake Lewisville v/ater. UV intensity = 0.375 W/L.
-------
0.2
D= I. 33mg/L - min
500 IOOO 1500
time, seconds
0.2
D = 0.775mg/L-min
500 IOOO I5OO
time, seconds
1.0
0.8
0.6
S/S.
2000
0.4
0.2
D= 0.48 mg/L-min
500 IOOO 1500
time, seconds
2OOO 2500
O
H1
1.0
0.6
0.6
S/S.
0.4
0.2
O
D =0.23mg/l_- min
SOO IOOO 15OO 2000 25OO
time, seconds
1.0
o.e
0.6
S/S,
0.4
0.2
D = 0. JO mg/L - min
500 1000 1500 2000 2500 3OOO
time, seconds
Figure A-ll. Tetrachloroethylene in Lake Lewisville water. UV intensity = 0.189 W/L.
-------
D= ].33mg/L - min
500 1000 1500
time, seconds
0.8
0.6
0.4
0.2
D = 0.775 mg/L - min
5OO
1000 1500
time, seconds
2000
1.0
0.8
0.6
0.4
02
D » 0.51 mg/L - min
500 1000 1500 2000
time, seconds
NJ
O
to
1.0
0.8
0.6
0.4
0.2
D= 0.295mg/L -min
500
1000 1500
time, seconds
2000
1.0
0.8
0.6
0.4
O
D * 0.11 mg/L- min
O
O o
O
500 1000 1500 2000 2500 3000 3500 4000 4SOO
time, seconds
Figure A-12. Tetrachloroethylene in Lake Lewisville water. UV intensity = 0.189 W/L.
-------
1.0
0.6
0.6
0.4
0.2
D= 1.33 mg/L-min
1.0
0.8
02
1000 2000 3000 4000
time, seconds
D = O.775mg/l_-min
i
i
1000 2000 3000
time, seconds
M
O
u>
i.o
Q8
Q6
04
0.2
D = 0.448 mg/L -min
1000 2000
time, seconds
3000
1.0
OB
0.6
0.4
0.2
04
0.2
1000 2000 3000
time, seconds
4000
D = 0.11 mg/L -min
1000 2000 3000
time, seconds
4000
Figure A-13. Chloroform in Lake Lewisville water. UV intensity = 0.375 W/L.
-------
1.0
0.8
06
0.4
0.2
D= 1.33 mg/L- min
1000 2000 3000
time, seconds
10
08
0.6
\
0.4
0.2
4000
D = 0.775 mg/L-min
looo 2000 3000 4000
time, seconds
to
o
1.0
0.8
06
0.4
0.2
Q D = 0.454mg/L - min
O
1000 2000 3000
time, seconds
1.0
Q8
oe
02
4000
D = 0.26mg/l_-min
1000 2000 3000
time, seconds
1.0,
Q8
0.6
0.4
Q2
D = 0.113 mg /L- min
4000 0 1OOO 2OOO 3000 4000
time , seconds
Figure A-14. Chloroform in Lake Lewisville water. UV intensity = 0.189 W/L.
-------
1.0
0= 0.30 mg/L min
10 20
30 40
time, min
90 60 70
1.0
D= 0.53 mg/ L min.
10 20 30 40
time, min
50 60 70
ISJ
O
1.0
D=0.80 mg/L min.
10 ZO
30 40
time, min
90 60 70
1.0
D = 1.33 mg/L min.
10 20 30 40
time, min
50 60 70
Figure A-15. Broraodichloromethane in Lake Lewisville water. UV intensity = 0.375 W/L.
-------
1.0
0.5
O 10 20
D= 0. 30 mg /L- min
O
30 4O 50 60
t, min
7O
1.0
0.5
0 10 20
D = 0.53 mg / L- min
30 4O 50 6O 70
t , min
NJ
O
1.0
0.5
D= 0.80 mg /L - min
1.0
C
0.5
0 10 20 30 40 50 60 TO
t, min
D = 1.33 mg/L min
0 10 20 30 4O 50 60 70
t, min
Figure A-16. Bromodichloromethane in Lake Lewisville water. UV intensity = 0.189 W/L.
-------
APPENDIX B
CHEMICAL KINETICS OF OZONE/ULTRAVIOLET-INDUCED REACTIONS OF
ORGANIC COMPOUNDS IN WATER - EXPERIMENTAL DETAILS
CONSTRUCTION OF THE REACTION SYSTEM
The ozone contacting reactor used in this work was a 6 inch
O.D. by 10 inch high quartz bell jar with flat flanges, which
was purchased from Quartz Scientific Inc., Fair Park Harbor,
Ohio (Figure 11, main report). The cover, impeller, sparger and
baffles are fabricated from Pyrex glass and fitted to the bell
jar by Mr. Tom Denton, glass-blower of the Department of Chem-
istry, North Texas State University. The impeller is driven by
an electric motor with glass shaft through a PFTE 0-ring. The
stirring speed is controlled by a variable transformer which is
calibrated using a RMP Tachometer Model C-891, Power Instruments,
Inc., Skokie, Illinois. The openings on the cover are used for
inlet and outlet of ozone/oxygen gas mixture, also for intro-
ducing substances to the reactor and withdrawing samples for
analysis. Stainless steel Swagelok nuts and Teflon ferrules are
used for sealing the openings. The whole reactor is set and
locked into an aluminum stand made by the machine shop of the
Physics Department, North Texas State University. The dimen-
sions of the entire reactor are such that it fits snugly into a
Rayonet Photochemical Reactor, the Southern New England Ultra-
violet Co., Middletown, Connecticut,Owhich is equipped with low
pressure mercury lamps, R.P.R. 2537 A, concentric to the reac-
tor axis. Up to £ total of sixteen lamps may be used, each of
which is rated at 14.0 watts. The spectral energy distribution
of the 2437 A lamp is shown in Figure B-l.
The geometry of the contacting reactor is designed such
that when three liters of solution is contained it will be fully
exposed to the radiation of the lamps.
Figure B-2 shows the schematic diagram of the photolytic
ozonation system. Oxygen (O) is fed to the generator, Grace
Model LG-2-L2, W. R. Grace & Co., Columbia, Maryland, via valve
(A) . The concentration of ozone in the gas can be varied either
by variation of gas flow (B) or primarily by the generator reac-
tion chamber pressure and voltage. Under certain conditions a
portion of the gas from the generator is vented to the ozone
destroyer (P) so as to give the desired flow of 03/02 gas into
the reaction flask.
20.7
-------
DC
UJ
UJ
DC
100
90
80
70
60
50
4O
30
20
10
SPECTRAL ENERGY DISTRIBUTION
2537A REGION
250
WAVELENGTH
300
MILLIMICRONS
350
INTENSITY READINGS
WAVE LENGTH (A)
2537
2652
2804
2894
2967
3022
3129
3654
4047
4359
5461
5780
IN MICROWATTS/cm AT.1
2 INCHES FROM LAMPS
16000
488
16
22
83
40
313
267
316
960
523
1 13
Figure B-l. Spectral distribution from ultraviolet lamp,
2:08
-------
Vent
fO
1 Generator
|
(A) VALVE
(B) ROTAMETER
(C) FLOW CONTROL VALVE
(D) ROTAMETER
(E) SIX PORT VALVE
IF) ROTAMETER
(6,H) OZONE MONITOR
(I) RECORDER
(J) PHOTOCHEMICAL REACTOR
(K) ENTRANCE TO REACTOR
(L) EXIT FROM REACTOR
(M,N) STIRRER AND CONTROLLER
(P) OZONE THERMAL DESTRUCTION UNIT
(0) VENT LINE
Figure B-2.
(a) Schematic diagram of the photolytic ozonation system.
flow for ozone off gas measurement.
(b) Alternative
-------
Flow of the gas stream in the reaction system is controlled
by a six-way valve (E), made of one disc of stainless steel and
one block of Teflon. The valve may be connected in such a way
to direct either inlet 0->/O2 (K) or the off-gas stream (L) to
the UV cell (G) of an ISCO Model UA-5 Absorbance Monitor (H),
Instrumentation Specialties Co, Lincoln, Nebraska (Figure B-2
insert). The monitor is originally designed for liquid chroma-
tography. In this work, 03/02 mixture is passed through the
sample cell and oxygen from the gas cylinder (O) is used as
reference. The readings of the monitor at 254 nra are calibrated
to indicate the ozone concentration in the mixture. In order to
monitor the changing of ozone concentration during the runs, an
OmniScribe recorder (I), Houston Instrument, Houston, Texas, is
connected to the UA-5.
In the normal flow configuration, the flow entering the
contacting reactor is controlled by valve (C) and the flow rates
entering (K) and exiting (L) from the reactor are measured by
Fisher & Porter rotameters (D) and (F) respectively.
The 03/O2 mixture exiting from the generator is split into
two streams so as to give the desired flow of Oj/O2 gas into the
reactor. One stream (Q) goes directly to the destroyer (P) and
the other goes to the six-way valve (E). The destroyer is main-
tained at 350°C. All the 03 in the split stream (Q) and reactor
off-gas stream (L) is destroyed by passing through the heated
tube and vented to an exhaust hood. The 0-3/02 mixture in the
reactor off-gas stream may be directed into the UA-5 monitor by
the six-way valve for ozone concentration measurement as shown
in the insert to Figure B-2.
The Rayonet Photochemical Reactor (J) has a fan on the
bottom of the inside chamber to vent the heat generated from the
lamps. The lamps are connected to the lamp-holders mounted on
the shiny aluminum wall.
MEASUREMENT OF OZONE DOSE RATE, OZONE CONCENTRATION IN SOLUTION,
AND ULTRAVIOLET RADIATION FLUX
Three liters of boric acid buffered potassium iodide solu-
tion (BKI) (150) made from deionized, carbon filtered water is
placed in the reactor and stirring begun. The gas flow rate of
the O3/02 mixture into the contacting reactor is measured, and
samples of the resulting iodine solution are withdrawn. The
concentration of iodine is determined colorimetrically at
A = 352 nm on a Spectronic 20 (Bausch & Lomb, Rochester, N.Y.),
which is previously calibrated by standard iodine solution (151).
By knowing the time between two samples and the flow rate, the
ozone dose rate in mgO^/min-L liquid can be calculated for a
fixed reading on the UA-5. A calibration curve of absorbance
vs. concentration of ozone in the gas mixture is constructed for
direct ozone concentration measurement with the UA-5. The
210
-------
stoichiometry of the ozone/iodine reaction is assumed to be
2H+ + 03 + 21- ^ i2 + Q2 + H20
Measurement of dissolved ozone in solution is as follows.
A portion of the solution to be analyzed (1-5 ml., depending on
the concentration of ozone) is withdrawn andmixed immediately
with BKI solution to make a total volume of 10 ml. The resulting
iodine absorption is measured on a Spectronic 20 and compared
with a calibration curve as described above. The concentration
of ozone in solution is expressed as mg/Lliq .
The effective radiation transfer rate from one or more
lamps of the Rayonet Photochemical Reactor is measured by fer-
rioxalate actinometry. The calculation of photons being trans-
ferred into the reactor is adapted from Parker (152). The
effect of gas flow and stirring rate on radiation transfer rate
is also measured.
MATERIALS
The water used in this work was prepared by purifying the
deionized, distilled water which was supplied by Ahlfinger's
Water Corp., Dallas, Texas. The water in a five-gallon glass
jar was forced by nitrogen through a four foot long by one inch
O.D- glass tubing containing 22 inches of Filtrasorb 400 acti-
vated carbon (Calgon Corp.), 12 inches of X&D-4/8 (Rohm and
Haas), 1 inch of Bondapak C18/Corasil (Waters Associates) and a
0.45 micron filter (Millipore Corp.). The TOG content of this
water was checked periodically with a Dohrmann DC-54 Organic
Carbon Analyzer and generally was below 600 yg/L. The resulting
water was used for all the kinetic studies in purified water
medium and for determination of by-products in this work. The
purified water for ozonation and photolytic ozonation runs was
further purified by applying a high ozone dose rate with UV
radiation for at least twenty minutes to destroy ozone demand
which might be present. Water temperature was ambient. For
kinetic studies in purified water, no pH adjustment was made, so
that no additional ions would be present to serve as radical
scavengers. The pH of this purified water was 6.5 - 7.0. When
a different pH was needed, it was adjusted using NaOH and H2S04
solution for pH = 4, 6 and 8, and Borax buffer for pH = 10.
Water was collected from Lake Lewisville, Texas, and stored
in five-gallon glass jars with good sealing. Before being used
as a medium for photolysis, ozonation and photolytic ozonation
the water was double-filtered through Whatman Quantitative 1 and
4 filter papers. The carbonate, bicarbonate, ammonia and COD
contents had been checked according to Standard Methods (157)
and TOC was also measured.
211
-------
Model compound HCB was either obtained from Analabs, Inc.,
North Haven, Connecticut or synthesized.in the laboratory by
Oilman's method (153-156). 2,4,6-trichloroaniline was obtained
from Aldrich Chemical Corporation, Inc.; sodium nitrite and
copper dust (electrolytic grade) from Fisher Scientific Co.;
pentane, methanol and hexane, nanograde, from Mallinckrodt, Inc.;
internal standard, hexachlorobenzene, from EPA, Research Triangle
Park, N.C. All other standard laboratory chemicals were of
reagent grade.
PREPARATION AND ANALYSIS OF SAMPLE SOLUTIONS
Stock solutions of chloroform, bromodichloromethane, and
tetrachloroethylene were prepared by stirring pure compound in
contact with water until it dissolved to produce an almost sat-
urated solution. The concentration of this solution was deter-
mined by gas chromatography from comparison with weighed stan-
dard solutions. The reaction mixture was prepared by spiking
3 L of water in the reactor with the appropriate amount of con-
centrate to yield a concentration of 100 yg/L. Direct aqueous
injection was used on all dilute aqueous solutions after deter-
mining the blank value of the water itself, which was always low
(0-2 yg/L). Direct aqueous injection gas chromatography was
performed on a 2 mm i-d. X 6' chromosorb 101 (100/120) column
at 160° - 200°C depending upon the compound. Detection was by
Tracor or Hewlett-Packard 6^Ni electron capture, the signal of
which was electronically integrated.
Trihalomethane precursors (TTHMFP) were analyzed by chlor-
ination of the samples at an applied dose of 50 mg/L chlorine
for three days. Previous studies had shown that for Cross Lake
water, three days was sufficient time to reach the plateau in
trihalomethane (THM) formation. At the end of three days, ex-
cess chlorine was quenched with sodium sulfite and THMs ana-
lyzed by liquid-liquid extraction with pentane. The samples
were buffered to pH 6.5 using phosphate prior to the chlorina-
tion, and the samples were stored in the dark at 26°C for the
three day reaction period. Gas chromatography of non-aqueous
solutions was on a 6' X2mmi.d. column of 12^ % Squalane on
Supelcoport (100/120) using electron capture detection.
Because of the limited solubility of hexachlorobiphenyl
(HCB), a different procedure was used for preparation of HCB
solutions. Model compound, HCB, was weighed to 0.1 mg and dis-
solved in methanol as a stock solution. An 18 yl portion of
this solution was spiked into 3 liters of water to give the 60
parts-per-billion (yg/L) concentration which was used for all
the kinetic runs. A portion of the solution was withdrawn from
the reactor at 1, 3, 5, 10, 15, 20, 25, 30, 35, 40, 45, 50, 60,
70, and 80 minutes after the runs began. In some cases, zero
time sample was also withdrawn, but this practice was of quite
limited value since a finite mixing time is required to achieve
212
-------
homogeneity. Each sample was placed in a twenty-five mL glass
vial containing 0.5 mL of ozone quenching reagent (120 mg total
of Na?S03). The sample solution was withdrawn from the reactor
by using a glass syringe and used to fill the vial up to an
engraved mark on the vial indicating a 20.5 mL volume. One mL
of hexanes containing internal standard, hexachlorobenzene, was
added. The vials were put on a Lab-Line Junior Orbit Shaker,
Lab-Line Instruments, Inc., Melrose Park, Illinois, and shaken
at the highest speed for twenty minutes. A 3 yL aliquot of the
organic layer was injected into the gas chromatograph for
detection of internal standard and HCB.
Gas chromatographic (GC) detection of HGB and internal
standard was performed on a Model 560 isothermal gas chromato-
graph, Tracer, Inc., Austin, Texas, equipped with Ni-63 electron
capture (EC) detector. A 1 foot by % inch O.D. glass column
containing 10% SP-2100 on 100/120 mesh Supelcoport, Supelco,
Inc., Bellefonte, Pennsylvania, was used. Carrier gas was 10%
methane and 90% Argon, Union Carbide Corp., Linde Division. The
oven temperature was kept at 176°C; injection port and detector
were at 200°C and 350°C respectively. A programmable computing
integrator, Supergrator-3, Columbia Scientific Industries,
Austin, Texas, was used to quantify and record the heights and
areas of the GC peaks.
PROCEDURE FOR KINETIC RUNS
During the kinetic runs, it was desirable insofar as pos-
sible to have the ozone flow and UV light equilibrated within
the reactor before substrate was added. Because of the required
mixing time, this procedure did not give a true zero-time
measurement of substrate concentration, but surmounted severe
mathematical difficulties involved with a non-constant ozone
concentration. In all cases where the reaction was first-order
in substrate (all model compounds but HCB), it was the slope of
the semi-log plot which was of importance, as will be explained
later, so that the exact initial concentration did not have to
be known. Therefore, the spiking of substrate was done in such
a way as to introduce a nominal 100 yg/L (+5-8%) of substrate
at time = 0 into a steady-state ozone/UV system. Furthermore,
since 10-30 ymoles/L of ozone per minute were being introduced
into a system to which 0.6 - 0.9 ymoles/L of substrate had been
added, perturbation from that steady-state was minimized, pro-
viding a good approximation to a pseudo-first order system.
This procedure avoided having to solve the simultaneous differ-
ential equations which would result if the reactor could not be
assumed to be in a steady state at the beginning of the reaction.
The experimental procedure, therefore, was to place three
liters of purified water into the reactor, then introduce the
desired ozone flow and dose rate into the reactor with the
desired intensity of UV present, until the ozone concentration
stabilized. The necessary amount of aqueous solution of test
213
-------
compound was then added and the timer started. Samples were
withdrawn periodically by syringe and directly injected into the
chromatograph for analysis. During the experiments, ozone con-
centration in the liquid and in the gas streams into and out of
the reactor was monitored. In the case of tetrachloroethylene,
the higher dose rates gave a reaction which was so fast that
only one data point could be collected before the substrate in
the reactor was depleted below the detectable level. In these
cases, substrate was respiked and the sample withdrawn after a
different elapsed time. This process was repeated until a suf-
ficient number of points had been collected to construct a dis-
appearance curve.
Since the solubility of HCB in water is very low, it was
first dissolved in methanol, as described previously. To begin
a run, 18 uL of that stock solution was spiked into the water in
the reactor, resulting in a substrate concentration of 60 pg/L.
Sampling and analysis then preceded as described in the previous
section.
Experiments to determine the rate of destruction of sub-
strate by ozone without UV (ozonolysis or "ozone-only") were
conducted identically but with no UV radiation present. Exper-
iments to determine the rate of destruction of substrate by UV
with no ozone present (photolysis or "UV only") were conducted
identically to the ozone/UV experiments but the 0-/O- gas stream
was replaced by nitrogen at the same flow rate. Experiments
with nitrogen flow and no UV were conducted to determine the
rate of purging of substrate from the reactor in the case of a
volatile substrate or of adsorption to the walls in the case of
a hydrophobic substrate such as HCB, and these rates used to
correct the experimental results of other runs.
When investigating the kinetics of destruction of the test
compounds in a natural water matrix, the ozone flow and UV radi-
ation could not be equilibrated prior to the beginning of the
experiment because of the degradation of the natural organics
which would occur. For. these experiments the natural water was
spiked with test compound and the preset ozone concentration and
UV turned on at time = 0. Sampling was then conducted as des-
cribed previously. As with purified water, a sample of the
natural water itself was anlyzed to provide a blank to be sub-
tracted from results obtained during the run. In almost all
cases, the blank was found to be negligible.
EXPERIMENTAL DIFFICULTIES
Since ozone is so reactive and the ozone/UV process in par-
ticular is not well understood, it is worthwhile to mention a
few of the experimental difficulties encountered in this inves-
tigation, in order to save other investigators from encountering
the same pitfalls.
214
-------
The single greatest difficulty encountered was the consis-
tency of the water used in the experiments. When a natural ma-
trix was used, 20-30 gallons was collected at one time since it
was found that even just a couple of weeks later, the lake water
might have greatly different properties. It was not practical
to store the entire 20-30 gallons at 4oc so series of experi-
ments were performed continously over a period of just a few
days. Even under these conditions it could be seen that the
water was changing somewhat during the period of storage. Natu-
rally this change is to be expected in an actual treatment situ-
ation also, but in a kinetic study of this type, small changes
can affect the apparent dependency of the reaction rate on, for
example, UV intensity more than the actual rate of any one ex-
periment appears to be affected. It is strongly recommended
that investigators store natural water at 4°C and perform groups
of experiments on the same water in the shortest possible time
span. The addition of chemical preservatives is questionable
without further investigation, due to the radical scavenging
properties of many substances.
A source of even greater exasperation at the beginning of
these studies was the difficulty of producing consistent good
quality purified water. Reaction rate for indiscriminate free
radical species with mixtures is shown in Section 4 of the main
report to be proportional to the quantity
kM[M] Q"1 = kM [M] {Sk^S^T1 ,
where [S.] is the concentration of substance i present and k. is.
its reaction rate constant, and M represents the substrate of
interest. Assuming for simplicity that the reaction rate con-
stants are equal and that only one other component, S, is
present
1 M
M+S •
When M is large and the impurity S is small
M/(M+S) a 1 .
But when M is small and S is large, as in the case of a micro-
pollutant in natural water,
M M < ,
M+S S
For the case of purified water, M « 100 yg/L while S = TOC
may be 100-600 yg/L. Clearly in this situation it is more im-
portant to have consistent water than extremely pure water,
215
-------
since compared to substrate level, the water cannot be abso-
lutely pure. A natural consequence of this situation is that
the purified water kinetics are probably not absolute rates, but
are to be taken as the practical upper limit—a useful reference
point—with the natural matrix studies at the other end of the
scale.
As an aid in obtaining pure water, the water used in the
purified water kinetic runs was "burned" with ozone/UV before
introduction of the substrate. The effect of this burning is
shown in Figure B-3, where two consecutive runs using the same
water were performed. The data shown in the figure is from work
early in the project, using water which was not as clean as that,
used in the kinetic runs reported below, and so the effect is
exaggerated. Presumably, organic matter is present intially so
that the substrate must compete with organic carbon for the
active species. As the interfering matter is removed, the rate
of substrate removal increases until the final rate in curve A
is seen to approach that in curve B, where the same experiment
has been rerun using the same water. In the later kinetic ex-
periments , it is not known how much organic matter remained
after the burn, hence the statement about these results being a
practical upper limit. Other issues relative to this point will
be discussed in a later section.
Another difficulty encountered at first was ozone measure-
ment. The UV absorbance procedure described above is strongly
recommended since it not only provides a continous monitor of
the incoming and outgoing ozone stream, but also because is does
not necessitate interruption or change in the back pressure of
the gas stream in order to perform a measurement. The strip
chart recorder informs the operator at a glance if an unnoticed
power surge has occurred, not an uncommon occurence in these
investigators' laboratory.
A third problem encountered in this investigation was the
tendency of the power output of the UV bulbs to change over the
weeks. It was found necessary to perform the actinometric cal-
ibration frequently to ensure a correct value for the UV inten-
sity. It is not sufficient to merely state the manufacturers
figures for power output, transmission coefficient of quartz,
etc., as a 10% error can change the final kinetic expression
considerably.
PROCEDURE FOR KINETIC RUNS ON DESTRUCTION OF TRIHALOMETHANE
FORMATION POTENTIAL
Water Sample Sources
Ohio River water was shipped to the location of the ozone
reactor (Houston Research Inc., Houston, Texas) by motor freight
in epoxy-lined 55 gallon drums. Water from Caddo Lake, Texas
216
-------
In
[CHCI3]
pH6.5 2.8 my 03 .42//UV
/W»lr* 3 /
B(A) PH6.5
'mm
'mm
UV
WATER REUSED FROM RUN A
"4 10 20 30 40 50 60 70 80 90 100
TIME(MIN.)
Figure B-3.
Effect of water purity on photolytic ozonolysis
of chloroform.
217
-------
was shipped in 5 gallon glass bottles by private automobile.
Ohio River water was used as received; lake water was run un-
filtered and after filtration through S & S No. 588 paper.
Humic acid (Aldrich Chemical) solutions were prepared in com-
mercial distilled water which had been filtered through acti-
vated carbon (Calgon Filtrasorb 400) and an XAD-2 polisher (Rohm
and Haas).
Ozonolysis and Qzone/UV Treatment
The reactor used in these studies has been described pre-
viously (103). It is a stainless steel cylindrical reactor of
approximately 22 liter volume, designed so as to achieve maxi-
mum gas-liquid transfer. Ozone is generated by a W. R. Grace
Model LG-2-L2 Corona generator using high purity oxygen. The
transfer lines from the generator are stainless steel, and pro-
vision is made to divert the 03/02 stream to an iodide trap for
calibration of the ozone content entering the reactor. In later
runs (on Caddo Lake water), provision was added to allow exit
gases from the reactor tq be trapped so that a calculation of
ozone consumed could be made. Ports in the reactor allow for
removal of samples during the course of a run, addition of rea-
gents to the reactor, temperature measurement, etc. Two quartz
wells extend into the reactor from the top; they house two
Hanovia high-pressure mercury-vapor lamps, at 200 and 450 watts,
respectively. A coil of metal tubing around the reactor may be
used for cooling during photochemical runs; however, there is
still some rise in temperature when lamps are used for sustained
periods. Oxygen flow rates of 200 ml/min (STP) were used in all
runs reported here. Ozone content of the gas stream leaving the
generator was calculated by the iodide-thiosulfate procedure,
and is expressed as mg03/l (STP) or as W/W%. Under the condi-
tions used, 1% W/W 03/02 corresponds to a dose rate of 14 mg
03/3! of gas or 2.8 mg03/min.
The procedure for making an ozone or ozone/UV run was as
follows: Twenty liters of river or lake water was charged into
the reactor, which had been "cleaned" with deionized water plus
ozone/UV for at least four hours. The reactor was capped and
ozone flow introduced at a flow rate of 0.2 sl/min with a pre-
calibrated ozone content. In the case of ozone/UV runs, either
one or both of the lamps was turned on approximately five min-
utes prior to the introduction of the ozone.
In the case of humic acid runs, 20 1 of "organic free"
water was charged into the reactor and ozone flow (20 mg03/min)
was initiated. After the ozone had stabilized at 2.2 mg/L, a
concentrated solution of humic acid in 0.1 N NaOH was added so
as to give a total humic acid concentration of 1 mg/L. Enough
6N sulfuric acid was immediately added to adjust the pH to 7.0-.
Ozone flow at 0.2 sL/min (7% W/W) was continued. In the case of
218
-------
runs, the lamps were turned on at the moment of humic acid
addition.
Sampling Procedure
Precursor levels were measured by chlorination of samples
withdrawn from the reactor at time zero (before ozonation) and
at intervals thereafter. Triplicate samples were withdrawn
directly into 120 mL septum bottles containing excess ozone
quenching agent (solid Na?S03) and phosphate buffer, so as to
give a final pH value of 6.5. The samples were immediately
topped with PTFE-lined septa and sealed with crimped aluminum
caps. These three'Samples were stored in ice and later shipped
by air freight to Denton for analysis in the NTSU laboratory.
One other sample drawn at the same time was used on-site to
measure pH and residual ozone by the iodide-thiosulfate
procedure.
Shipping time to Denton was approximately five hours in
most cases. Upon receipt, the chilled samples were stored im-
mediately in a standard refrigerator and processed after no
longer than a sixteen hour total delay since sampling.
Chlorination of all samples was carried out by the same
procedure. Chlorine water was added to each bottle with a
microliter syringe through the septum. Sufficient chlorine was
added so as to give the desired residual, usually 10 mg/L for
Ohio River samples and 15 mg/L for Caddo Lake samples.
Samples were stored in the dark at 25+1° for seven days.
On the final day, trihalomethane analysis was carried out by
the pentane liquid-liquid extraction procedure developed in
this laboratory (167). Generally, chlorine residual at the
final day was in the 1.0-2.0 mg/L range.
219
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APPENDIX C
CHEMICAL KINETICS OF OZONE/ULTRAVIOLET-INDUCED REACTIONS OF
ORGANIC COMPOUNDS IN WATER - DETAILS OF KINETIC ANALYSIS
FOR FIRST-ORDER SUBSTRATES
A MATHEMATICAL TECHNIQUE
The following is a brief explanation of a useful mathe-
matical technique which is used extensively throughout the
following sections. This explanation is intended to make the'
following section on kinetics more readable to those with
limited mathematical background, and no attempt is made to be
rigorous.
It is a property of logarithms that the log of a product
is the sum of the logs:
In AB = In A + In B.
Another property of logrithms is that the log of a quantity to
some power, b, is equal to b multiplied by the log of the
quantity.
In Bb = b In B.
/
Combining these two properties gives the following expression:
In ABb = In A + b In B.
This expression is useful in kinetic analysis since rate laws
frequently have the form
- || = Rate = kSa .
Taking the log of both sides yields
In Rate = In k + a In S, [C-l]
which is the equation of a straight line. Therefore, if In Rate
is plotted vs.In S, the line will have slope a and intercept
In k. Use of this technique allows the step-by-step evaluation
of the dependency of a rate expression on various parameters,
as shown in several of the following sections.
220
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KINETIC ANALYSIS OF OZONE/UV REACTIONS WHICH ARE FIRST ORDER IN
SUBSTRATE
I. Purified Water as the Reaction Medium
AI Tetrachloroethylene—Since the rate of purging was
found to be insignificant compared to the reaction rate of TCE,
the purging term in eq. 71 of the main report was omitted and
the data was modeled using a three-term equation of the following
form:
_j a b .
(ozonolysis) (photolysis) (0.,/UV)
Where S = substrate concentration in moles-iT1
I = UV intensity in W-L"1
(03) 1 = ozone concentration in liquid, in moles'L~
D = ozone dose rate, in mg-L~ min , where L is
liters of substrate solution.
-In addition, k , k , k = appropriate rate constants, and
a, b, c, d, i, m, n = reaction orders, all to be evaluated
experimentally.
The experiments using ozone alone were analyzed by first
plotting In S/S vs. t. Within the precision of the data, these
plots were linear, indicating that reaction could be considered
first order in substrate (a=l). The experiments are carried out
by equilibrating ozone flow with purified water, then spiking
lOOyg/L of substrate (water concentrate) into the reactor at
t=0, in order to have (O3)-L constant. Since there is a finite
mixing time, there are temporary inhomogeneities in the reactor
and the reaction has slightly different characteristics at the
very beginning. Therefore, the following procedure was used to
obtain rate constants.
b
For each experimental point, -(In S/So)/t = ko (03)-^ was
calculated and tabulated, as below:
t,sec: 390 730 1040 1400 1740 2160 2495 2845 3240
105 x (-In S/S )/t: 153 121 115 107 109 109 106 105 105
Clearly, after the first couple of points, the reaction settles
into first order. The procedure used was to start at the end of
the table and work toward the front calculating the mean and
standard deviation, then applying the 4a rule (160) for outlying
221
-------
results to each successive value of (-In S/SQ)/t until some value
failed to fall within four standard deviations (a) of the
previously calculated mean. For the example given, only the
first value was eliminated. The logarithm of the resultant K
(average) was plotted along with error bars representing a for
the log of each ozone dose rate. This plot is shown in Figure
C-l. The dotted line is a least squares fit of all five points,
and is seen to miss the error bars on three of those five points.
This line has a slope of 0.76. When the outlying point is
eliminated, linear regression yields a line which fits all
remaining points and has a slope of 0.98. This line was
selected, the exponent forced to unity, and the appropriate
correction made in the rate constant, yielding
— f^ Q
~-^j7 = k (0_),S (for ozonolysis)
where k = 5.87 L mole sec
o
To evaluate the photolysis term,
-*§ _ k ,c sd
dt ~ kp J S
fi
k I was taken as the slope of the (linear, if first order in S;
i?e. , d=l) In S/S vs. t plot, and was calculated for 1, 2 and 4
lamps for two different dates (and therefore two different UV
intensities, due to lamp deterioration). Of these data, the old
1 lamp data was found not to be first order in substrate and was
eliminated. The remainder of the kp 1° values were divided by
their experimentally determined I values to determine if they
were first order in I (c=l). Again, the discarding of one
outlying value (4 a rule) gave
-5-ir = k IS (for photolysis)
where k = 9.46 + 0.58 x 10~3 L-w"1 sec"1.
In order to evaluate the ozone/UV term of the rate
expression, In S/SQ vs. t plots were made of the ozone/UV data,
and were found to be linear; therefore
-(In S/SQ)/t = K = {ko(03)1 + k I + ku!iDm} .
The first two terms in K were calculated from experimental data
and the previously determined rate constants, and subtracted
222
-------
-8.0
-75
N)
-7.0
-6.5
-9.6
-9.4 -9.2
-9.0
In C03],
-8.8 -8.6
moles/ I.
-8.4 -82
-8.0
Figure C-l. Determination of order and rate constant of ozonolysis of tetrachloro-
ethylene.
-------
from the experimental values of -(In S/So)/t for the ozone/UV
experiments. The log of the. resulting 3 = k^I^-D was plotted
vs. In D, i.e., In 6 = In knl1 + m In D as before, to give
Figure C-2. Linear regression gave the top two curves for 2
and 4 lamps, and the dotted line for the one lamp plot.
Reconsideration of the raw data suggested elimination of the
experiment which used the highest dose rate-of ozone, and a
least squares fit of the remaining points yielded the solid line
shown drawn through the 1 lamp data. The resulting slopes are
0.96, 0.96, 0.93 for 1, 2, and 4 lamps respectively. Plotting
the intercepts vs. In I (Figure C-3) yields a slope of 1.04 and
a rate constant of 0.109. Forcing both exponents to unity gives
the final rate expression:
• _ -dS
'S ~ dt
V°3>1S
k IS
P
mole L sec~ 5.87+0.31
k IDS
u
-3
(9.46+0.58)xlO 0.109
L-sec mole
-1-1
L-sec W
2 -l -a -l
L mm sec mg w
At this point destruction curves were calculated to verify
the adequacy of the derived rate expression and the assumptions
used. These plots are shown in Figures A-1,2,3 (Appendix A)
where the equation is seen to describe the experimental results
within the limits of precision of the data.
B. Chloroform—In this system, the ozone/UV reaction is so
fast compared to purging, photolysis, and the ozone-only reaction
that the data can be well represented by just the ozone/UV term
of the rate equation,
"If = k iaDbSC = K SC = K'DbSC
dt u u u
As before, a plot of In S/SO vs. t was linear, indicating
that the reaction was first order in substrate (c=l) . For a
given UV intensity, I, plots of In Ku ys. In D gave b = 1.5 and
In Ku as the intercept. Plots of In K^(I) vs. In I gave
a = 0.266 and ku = 7.97X10"3 when time is given in seconds.
The final expression is
-dS _ 0.266 1.50 Q
dt ~ u D S
where
-1 • 50
224
-------
-7.0
Determination of substrate order in 0-/UV term
(tetrachloroethylene).
225
-------
In kur
-48
-46
-.44
-42
-40
n = 1.04
b = -2.213
-.38
-.36
-.34
-.32
-1.0
Inl
-2.0
-3.0
Figure C-3,
Determination of UV intensity exponent and rate
constant for 03/UV term (tetrachloroethylene).
226
-------
and the units of I and D are Watts/L and mg/L-minute, respec-
tively. The experimentally determined disappearance curves,
along with the curve fit calculated from the above equation are
shown in figures A-4,5,6.
C. Bromodichloromethane— An anomaly was found in the
results for bromodichloromethane (BDM) in purified water.
Unlike the two compounds discussed previously, for a given ozone
dose rate, runs using four lamps were found to be slower than
those using only two lamps, which were in some cases slower than
those using_only one lamp. Many of the experiments which
produced this unusual data were rerun and the results were found
to be exactly reproducible. The photolysis results were found
to be in the expected order, that is, with four lamps giving
faster destruction than with two lamps, and the photolysis was
found to be first order in UV intensity. Photolysis experiments
run months apart by different workers gave results so close to
each other that the data sets were combined to evaluate the
photolysis term (Figure C-4). As before, the ozone/UV pseudo-
first-order (PFO) rate constants were evaluated from the slopes
of In S/S0 vs. t plots (Figures A-7,8,9. The combined purging/
photolysis rate constants were subtracted from the total ozone/
UV PFO rate constants to separate out the portion due to the
ozone/UV reaction(s). Since ^^^a was below the detectable
limit, the ozonolysis term was neglected. When these PFO rate
constants, Ku*ff,D) were plotted against the ozone dose rate, D,
the curves were found to consist of two portions, a ramp and a
plateau, as seen in Figure C-5.
Inspection of these plots shows that as ozone dose rate is
increased from a value of zero, the rate of the ozone/UV
portion of the reaction increases linearly, until a cut-off
point is reached, after which increasing the dose rate no
longer increases the reaction rate. Furthermore, when the UV
intensity is approximately doubled in going from 0.084 W/L to
0.189 W/L, the dose rate cut-off approximately doubles also.
Apparently, the cut-off for 0.375 W/L was outside the range of
experimentation and would by inference be expected to lie at
slightly over D=2 mg/L min. The implication that there is a
relationship between UV intensity and maximum effective ozone
dose rate is of great importance from both a mechanistic and
practical point of view, and will be discussed in a later
section.
It is not known whether the non-coincidence of the three
y-intercepts in Figure C-5 is meaningful or a result of
experimental error. In this treatment it is handled as
experimental error, the 0.189 W/L intercept set equal to the
other two intercepts and the UV intensity dependence of the
rate constant determined as follows.
227
-------
2.5
2.0
1.5
C\J
1.0
O
0.5
TWO DIFFERENT WORKERS
O
D
.30
Figure C-4.
.10 .20
UV INTENSITY, W/L
Determination of photolysis rate constant for
bromodichloromethane in purified water.
228
.40
-------
c
E
CM
O
O 0.375 W/L
Q 0.189 W/L
A 0.084 W/L
0.5 1.0
OZONE DOSE RATE, mg/L-min
1.5
Figure C-5. Pseudo-first-order rate constants for bromodichloro-
methane in purified water.
229
-------
The total PFO rate constant K had the purging and photo-
lysis rate constants, kpurge and Kp, subtracted to give the PFO
ozone/UV rate constant Ku=kuIeDf (see equation 71,main report).
Plotting K vs. D gave a slope which represents 3 = kule.
Plotting In g vs. In I gives intercept ku and slope e (see
equation C-l and discussion). Linear regression of the points
comprising this plot gave ku = 6.56 X 10~2min~1(W/L)°•46°
(mg03/L min)"1 and the reaction order, e, of the UV intensity
equal to -0.466. The non-zero intercept of Ku vs. D shows that
Ku is actually of the form Ku = kc + ku!-°-46&D, so that the
total rate constant is
K = k +K+k+K
purge p c u
= k + k I + k + k i-°-466D
purge p c u
where kure=1.43X 10~3 min"1
k = 5.62 X 10~3 min"1 (W/L)"1
P
k = 8.78 X 10~3 min"1
u
= 6.56 X 10~2 min
(mg03/L min)
This equation should not be used for dose rates below the
lowest experimental value (0.15 mg/L min) because of the term
kc which is neither a function of UV intensity nor ozone dose
rate and yet is not present in the photolysis/purging experiment.
The term k undoubtedly arises from the oversimplification of a
very complicated kinetic situation.
These somewhat unusual results are discussed further in the
summary section which follows the rest of the kinetic analysis
details. The experimental data are shown in Figures A-7,8,9.
II. Lake Water as the Reaction Medium
A. Tetrachloroethylene—Destruction of tetrachloroethylene
in lake water by ozone/UV is sufficiently slower than in pure
water, that a term describing purging of TCE from the reactor
by the gas stream must be included in the rate equation.
The form chosen was
230
-------
dt
= (k
purge
(purging)
K
K
p
V
(ozonolysis) (photolysis) (O/UV)
The^ozone/UV reaction, as well as purging, ozonolysis, and
photolysis were all within experimental error of being first
order in substrate. As before, I represents UV intensity and D
is the ozone dose rate. Since the ozone concentration in the
liquid was continually changing throughout the reaction, the
rate equation was written in terms of ozone dose rate rather than
ozone concentration in the liquid. During ozone/UV runs, in
which ozone is destroyed by ultraviolet radiation, the concen-
tration of ozone in the liquid was monitored and the effective
dose rate (dose rate which gave the same profile) used in the
kinetic analysis, since in a strictly ozone reaction, only the
bulk liquid concentration is important.
Plots of In C/C0 vs. t for the ozone (only) reaction were
made and found to be linear. The corresponding rate equation is
p.
dt
= (k +
v purge
Da) S,
so that
I
S
dt
d In S
dt
purge
k D
o
is the slope of the above plots. These slopes were plotted vs.
D and the result was a straight line whose slope was ko and
whose intercept was kDurae- The fact that the plot was linear
indicates that a = 1. This plot is shown in figure C-6.
231
-------
o
O>
CO
1.0
ro "
2 0.8
x
o
X.
0.6
0.4
0.2
slope = 6.127 X I0"4sec"'
= k
intercept = I
655 X IO'4 sec'
k purging
05 1.0
OZONE DOSE RATE, mg/L-min
mg
L min
1.5
Figure C-6. Ozonolysis of tetrachloroethylene in lake water.
For photolysis experiments, the rate expression is
- It - "Wge + Vb> S = KPS"
The values previously obtained for kpurge were subtracted from
Kp to obtain numerical values of k I" for different I. Then
ln(k lb) = Iri k + b In I
•L
yields k and b when ln(k I ) is plotted vs. In I.
The rate constants in the ozone/UV experiments were
obtained by taking -(In C/Co)t~l as described in the previous
section on tetrachloroethylene in pure water. From these values
23.2
-------
of K were subtracted kp , K , and Ko (using D effective to
calculate Ko). The logarithm of the value of Ku thus obtained
was plotted vs. InD to determine the order of the factor D. It
was found to be 1.00+.09 for the 0.084 and 0.189 W/L experiments
and 0.469 for the 0.375 W/L experiments. Since the uncertainty
in the determination of K was almost twice as large for the
.375 W/L experiments as in the other two sets, and because in
the absence of data to the contrary a first order situation is
most likely, a value of d = 1.0 was used in subsequent calcu-
lations, subject to final comparison of calculated rate data
with that collected experimentally.
Intercepts of the previous plots gave values of In kula,
which, when plotted vs In I as before yielded a as the slope and
In ku as the intercept.
The final values of the constants in the rate equation
- |S + + l0.670 + ,-0.419 s
dt purge o eff p u
are:
k = 1.66 x 10~4 sec"1
purge
k = 1.31 x 10~3 sec"1(W/L)~°-67°
P
kQ = 6.13 x 10~ sec" (mg/L-min)~
11 —1 • —0 419
k = 2.97 x 10" J sec 1(mg/L-min) x(W/L)u'
u
For typical values of UV intensity I = 0.375 W/L and ozone
dose rate D = 1.33 mg/L min, the various terms in the rate
constant are:
= 1<66 x 10~ + 1<31 x 10~(-518) + 6.13 x 10
+ 2.97 x 10~3(1.33) 0.663
= 1.66 x 10~4 + 6.79 x 10~4 + 0 + 2.62 x 10~3
(purging) (photolysis) (ozonolysis) (ozone/UV)
It is seen that purging and photolysis contribute a greater
proportion of the total rate expression than in the case of TCE
in pure water, but that the dominant term is still the ozone/UV
reaction. In the case which has the highest Deff of those
233
-------
studied, I = 0.084 and D = 1.33, Deff =0.13 mg/L min, the
terms are
!_ dS = 1.66 x 10~4+ 2.49 X 10~4 + 7.97 x 105 + 1.40 x 103
S dt (purging) (photolysis) (ozonolysis) (ozone/UV)'
and ozonolysis still plays a minor role.
The integrated form of the above rate equation was used to
calculate disappearance curves for each of the ozone/UV experi-
ments performed, and the results are shown in figures A-10 -
A-12. Agreement with experiment is seen to be very good,
considering the opportunities for error accumulation in a
stepwise analysis of the type required here.
B. Chloroform-—Kinetic analysis on chloroform was carried
out using the methods described in previous sections. PFO rate
constants were determined for photolysis at two different UV
intensities and a plot of Kp vs I extrapolated to I = 0 to
determine the purging rate constant, kpurge. Ozonolysis was
performed at an ozone dose rate of D = 1.33 mg/L min, and the
purging rate constant previously determined, was subtracted from
the slope of a plot of S/S vs t for the ozonolysis. The
resultant PPO rate constant for ozonolysis was divided by the
ozone dose rate to give a second order rate constant, ko = 4.26
X 10~5 sec~l(mg/L min)~ , where
OZONOLYSIS RATE = ~||- = k D S
dt o
The concept of dose rate was introduced because even though
the presence of UV generally kept the aqueous ozone concen-
tration below the detectable limit, a relationship to the amount
of active specie generated was needed. When trying to calculate
the extent of ozonolysis which occurs in the presence of UV in
purified water experiments, the effective dose rate was
defined as that dose rate which in a non-UV experiment resulted
in the same equilibrium ozone concentration as was found in the
experiment, and was thus generally zero. In the case of
natural water experiments, the ozone and UV could not be
equilibrated before adding substrate because of the resultant
changes in the matrix and the effect of those changes on the
reaction rate. Therefore, in natural matrix experiments the
effective dose rate is by analogy defined as the dose rate which,
in an ozonolysis (non-UV) experiment produces the same aqueous
ozone concentration profile as seen in the ozone/UV experiment
being analyzed. In the chloroform experiments described here,
the effective dose rate was found to be 50-80 times smaller than
the actual dose rate, with the result that the ozonolysis itself
contributes very little to the disappearance rate of substrate."
234
-------
As in previous examples, the ozonolysis and photolysis
contributions were subtracted out of the total pseudo-first-
order rate constants for the ozone/UV experiments to yield the
PFO rate constant for the ozone/UV portion of the reaction,
K = k Tanb
*u Ku1 D
For each set of experimental conditions, In Ku was plotted
vs In D to yield intercept In kula and slope b = 0.562 + 0.212.
The values of the intercepts of the two lines representing the
two different UV intensities were so close together that only a
very slight UV dependence was seen (Figure 14, main report),
resulting in the following expression.
v - v -r-0.028 n0.562
u ~ *u I D
* k D°-562
U
where KU = 1.53 X 10~4 sec'1 (mg/L min)"0'562
= 9.18 X 10~3 min"1 (mg/L min)"0'562
The total rate expression which results is
_S = -||= k + k Deff + k I + k D°'562
dt purge o eff p u
= purging + ozonolysis + photolysis + CK/UV
where kpurge = 1.59 X 10"5 sec'1
k = 4.26 X 10~5 sec"1 (mg/L min)"1
o
k = 7.63 X 10~5 sec"1 (W/L)"1
P
k = 1.53 X 10~4 sec"1 (mg/L min)"0'562
The experimental disappearance curves, along with the curve fit
calculated from the equation 86 are shown in Figures A-13 and
A-14.
C. Bromodichloromethane—As in previous data sets, the
first approach taken was to analyze the ozone-only and UV-only
data in order to subtract their contributions from the experi-
mental ozone/UV rate constant. Although a few of the experi-
mental curves have a slight "less than first order" look
235
-------
(the semi-log plot curves slightly downward), it is small enough
that the data can be adequately represented using an expression
which is first order in substrate. This was verified by com-
parison of the calculated disappearance curves with those
obtained experimentally.
Because of time limitations, only two different ozone
concentrations were run in the ozone-only experiments. Similarly
only two different UV intensities were used in the UV only
experiments, so that one cannot absolutely distinguish between
various orders for these parameters. The procedure used was to
assume that these reactions were first order with respect to
ozone and UV respectively, and that substrate disappearance was
due to reaction plus purging from solution. The pseudo-first-
order rate constant for the overall disappearance was plotted
vs. the appropriate parameter (ozone dose rate, D, or UV
intensity, I) and the straight line defined by the two points
extrapolated to the y-axis, where D or I is equal to zero.
Since this value should represent the rate constant for
disappearance of substrate in the absence of reaction, it was
taken as the purging rate constant. Values were compared from
the ozone-only and UV-only plots, as a check of the validity of
the above assumption of first order, and the values were found
to be 3.01 X 10~4 and 2.30 X 10~4 min"1 respectively - well
within experimental error.
During the ozone/UV experiments, the ozone level in the
liquid was found to be negligible. Therefore, to determine the
portion of the ozone/UV rate constant which is due to the
ozone/UV reaction, just the purging and UV-only terms were
subtracted, that is, the rate expression was assumed to have
the form
•Mr = KS - ZK S = (K + K ,. + K ) S
dt j_ i purge phot u
= (k + k I + k IaDb) S
purge p u
Allowing b to take on any value resulted in a dependency of b
on I. Even when b was fixed at b = 1, a satisfactory fit could
not be obtained, and it was concluded that, as in the case of
bromodichloromethane in purified water, the situation was more
complicated than strict adherence to the model allowed. The
idea of identifying a particular term with a particular reaction
was abandoned, and the data was empirically fit to the form
K = k + kI + kIaD .
236
-------
The resulting equation was
Kdnin"1) = -7.386 X 10~3 + 8.948 X 10~2 I + 5.038 X 10~3l"°<481
D,
The experimentally determined rate equations for the
corresponding ozone-only and UV-only reactions were
/") G A
= 3.01 X 10~4 + 4.86 X 10~3 D
dt
(purge)
(ozonolysis)
and
- §1 = 2.30 X 10~4 + 2.99 X 10 2 I
dt
(purge)
(photolysis)
respectively, where t is measured in minutes. For comparison,
selected values of the rate constants (in min"1) of the
ozonolysis, photolysis, and ozone/UV reactions are shown below.
D,mg/l minX^
0
.30
1.33
0
2.6 X
1.76 X
6.77 X
io-4
lO-3
10-3
.189
5.43 X
1.29 X
2.46 X
.375
io-3
10-2
io-2
1.10 x
2.86 X
3.70 X
io-2
io-2
io-2
The experimental disappearance curves, along with the
curve fit, calculated from equation 88 are shown in Figures
A-15 and A-16.
237
-------
APPENDIX D
RESULTS OF KINETIC ANALYSIS OF OZONE AND OZONE/UV
DESTRUCTION OF HEXACHLOROBIPHENYL IN WATER
The same general form for the rate expression is used for
HCB as for the previously discussed substrates, except that the
purging term is not applicable because of the low volatility of
HCB, which is a solid at room temperature. Studies were made
using the usual gas flow through the reactor but with no ozone
or UV present, to determine that HCB was neither stripped from
solution by the gas nor adsorbed onto the walls of the vessel.
Thus, the ozonation rate expression was of the form
- S = - dS = k0 Dm Sn [D-l]
dt
while the ozone/UV reaction was described by two terms,
- S = - dS_ = Ku IaD^Sc + k Idse . [D-2 ]
dt P
As before, the ozone concentration in solution was very low
during an ozone/UV reaction (although not, in general, zero
because of the higher dose rates used) and the ozonolysis term
was not included, any occurring ozonolysis being lumped into the
much larger ozone/UV term. The results are summarized below, by
reaction medium.
I. LAKE WATER AS THE REACTION MEDIUM
The ozone concentration profile in the ozonation of HCB in
water is shown in Figure D-l. It should be noted that the dose
rates used here are for the most part higher than those used for
the three compounds previously described. The ozone-oxygen gas
flow rate was kept at 0.20 L/min and the stirring rate at
500 rpm. HCB was spiked and well-mixed before the introduction
of ozone. The profile shows there was a short induction period
in only one case, following which the ozone concentration
increased smoothly. With UV radiation applied, a quasi-equili-
brium state resulted at first as shown in Figure D-2. During
this time it appears that an immediate ozone demand of the
matrix is being filled, as judged by the proportionate lengths
of the induction periods in Figure D-2 , and by the approximate
correspondence between the equilibrium levels of curves D = 1.16
233
-------
and D = 3.95 in Figures D-l and D-2 respectively and the lengths
of their induction periods. This ozone demand should have
little or no effect on the ozone/UV process, however, since the
great difference between the ozone levels in Figures D-l and D-2
indicate that most of the ozone accumulated in Figure D-l is
photolyzed in Figure D-2.
Based on equations in reference 145, one can obtain the r\
and KLa values from the ozone concentration profiles with or
without UV. The quantity n is the ratio of the equilibrium
ozone concentration predicted by Henry's law to that actually
found. Since KLa should not be reaction dependent when most of
the fast reacting organics are depleted, it was used to calcu-
late KD'S for ozone concentration profiles at various UV radia-
tion transfer rates. As discussed in reference 145, K is the
first-order rate constant for the disappearance of ozone. In
this work, the concept has been expanded to include photodecom-
position as well as autodecomposition. This will be explained
in greater detail in the section on mass transfer. The results
are shown in Table D-l. The values of n and K were propor-
tional to the UV radiation transfer rates as expected. The
destruction of HCB by ozonation is very slow as shown in Figure
D-3. As in the case of photolysis, ozonation is faster at the
beginning of the reaction and slows down rapidly. A simple rate
expression cannot be used to describe the whole reaction path;
however, the initial rates of ozonation of HCB in natural water
were obtained and are shown in Table D-2. Figure D-4 shows that
the rates are first order with respect to the ozone dose rate.
Ozonation rates at longer reaction times are significantly re-
tarded as compared to initial rates and appear to be nearly
independent of ozone dose rates. As shown in Figure D-l, the
concentration of ozone in the liquid phase is very high in this
region and has reached a steady state level dependent on the
dose rate. It is clear that ozonation in this region is rate
controlling, i.e. the rate is independent on the ozone dose rate.
Based on equation D-l the initial rate expression for
ozonation of HCB in lake water can be written as
- aM = KQD - aM ,
where a = CM/C°
-2 -1
with K = 1.38 x 10 L (mg ozone) at pH 8. In the limiting
region°the rate is independent of ozone dose rate.
Photolytic ozonation of HCB in lake water was carried out
at various ozone and UV dose rates, and at different pH values.
As an example, the photolytic ozonation data of HCB under
0.25 W/L and various ozone rates are shown in Figure D-5. The
destruction curves are smooth and the reaction is generally very
239
-------
fast compared to ozonation. The photolysis curve under the same
radiation rate is also shown for comparison. At long reaction
times it is significant that the reaction rate is approximately
equal to that of photolysis. Under these conditions it is pos-
sible that the photolytic ozonation process is retarded by radi-
cal scavengers produced in the reaction and only the direct pho-
tolysis process remains. The reaction order with respect to HCB
and dose rate in equation D-2 was determined by the following
procedure. A smooth line was drawn through the data points, and
the tangents to the curve at certain values of a were taken as
the values of -a . For each ft , the values were corrected for
photolysis by the equation of 03/UV rate = r = -d - fi , .
Equation [52] of reference (161; is used for the photoTysis in
this system.
£> . = k I1/2S
M,A p
-2 1/2 1/2
where k = 1.02 x 10 L ' /W ' min in purified water
= 1.36 x 10~2 L1//2/W1//2 min in lake water.
This is the equation which describes the behavior over the
greatest portion of the photolysis disappearance curves.
Values of r were fit to the following rate equation.
r = -a,. -;
M
,, , _i^m n
\. . = k I D a.,
M, A u M
[D-3]
From the log-log plots of r and CL., the slope yields the reac-
tion order n with respect to aM. The values of k I^D111 then were
calculated and used for obtaining m and i. The values of n, m
and i estimated from various runs are
n = 2
m = 1
i = 1
The values of ku's of various runs are listed in Table D-3.
Thus, the empirical equation for photolytic ozonation of HCB in
Lake Lewisville water is expressed as
-a., = k
M u
a
'M
+ k
.1/2
a
'M
where
= 1.31 x 10'1 LVW-mg
-2 TJs,
k = 1.36 x 10
P
-mn.
240
-------
This expression was solved for time, t, as function of aM
and the values of t vs. aM in each case were calculated by using
the HP 2000 computer system. As shown by the solid lines in
Figures D-5 and D-6, the agreement between calculated and exper-
imental points is satisfactory.
In order to make the comparison of ozonation and photolytic
ozonation of HCB, Figure D-6 shows the runs with the same ozone
dose rate at various UV radiation rates. The effect of UV radi-
ation on ozonation is obvious.
II. PURIFIED WATER AS THE REACTION MEDIUiM
In the kinetic study of ozonation and photolytic ozonation
in purified water, substrate is spiked into the water which has
been equilibrated with either ozone or ozone/UV. Under these
conditions, the concentration of ozone in the liquid phase is
constant.
Expressions for the rate of disappearance of substrate by
ozonation and photolytic ozonation are similar to those used in
natural water matrix (equations D-l and D-2 respectively).
In purified water, pH effects on the ozonation and photo-
lytic ozonation were studied in addition to effects of ozone
and UV dose. In the experiments, pH was adjusted prior to the
addition of substrate. The pH was not further controlled for
the values of 4, 6 and 8; however, borax buffer was used for
pH = 10.
Figure D-7 shows the effect of pH on the ozone degradation
rate of HCB in purified water. As expected, ozonation at
elevated pH is accelerated (34) presumably by the formation of
hydroxyl radicals, however, the effect is not dramatic. The
rate law of ozonation of HCB in purified water is determined
to be
-aH = k0 • D'2 • ajj
and the rate constants are listed in Table D-4,
expression for photolytic ozonation is
[D-4]
The rate
*M = ku D
a
'M
+ k
a
'M
[D-5]
where k = 1.02 x 10
~2
and values of ku are listed in
Table D-4. The pH effects on the 03/UV destruction rate of HCB
are also noticeable as shown in Figure D-8. Figure D-9 shows
the photolytic ozonation runs conducted in purified water at
pH = 8. Figures D-10 and D-7 compare kinetics in lake water and
241
-------
purified water at the same pH and similar UV radiation transfer
rate, from which one can see that the lake water matrix does not
interfere significantly with the destruction rate; however, the
effect on the "ozone only" reaction is more apparent.
Summary—The ozone and ozone/UV reactions of HCB have been
found to be kinetically complex and to slow down tremendously in
the region beyond about S/SO = 0.5. For this reason they have
not been included in the process design calculations, although
their kinetic analysis has been outlined above, omitting the
exhaustive mathematical details. Possibly because of the meth-
anol present, ozone/UV reaction of HCB in purified water was
found to be approximately the same speed as in lake water,
although the mathematical form of the empirical rate equation
differed.
TABLE D-l. VALUES OF n, KLa AND KD FOR OZONE PROFILES WITH AND
WITHOUT UV IN LAKE LEWISVILLE WATER
Runs
Ia D (mg/
(W/L) min-L)
Ka
KD*
NLLW-1 1
2 2
3 2
0
4 4
6 1
7 .12 2
8 3
12 2
13 .25 3
14 3
17 2
18 .49 3
19 3
.16
.19
.95
2.83 + .87
.04
.21
.22 8.93 + 1.47
.04
.16
.04 17.90 + 1.47
.95
.16
.03 34.88 + 9.77
.91
u
(1.78 +
X
(1.78 +
X
(1.78 +
X
(1.78 +
X
.32)
10 2
.32)
10-2
.321
ID"2
.321
10~^
(3.29 +
X
(1.41 +
X
(3.00 +
X
(6.01 +
X
1.65)
10"2
.26)
10-1
.26)
10-1
1.73)
10-1
* Average value from individual runs. From reference 145 and
equation 93 of the body of this report.
= 1 + M . In this work n= 1 + kd + kp = 1 + Kd
KLa
KLa
KLa
242
-------
TABLE D-2. INITIAL RATE, R , AND INITIAL RATE CONSTANT, Ko,
POT? O ^f'YM'A n"
(W/L) (mg/min-L)
1.21
.12 2.22
3.04
1.14
2.16
.25 3.04
3.95
1.21
.49 2.16
3.03
3.91
ku
(L2/W-mg)
0.132
0.125
0.108
(a) [HCBl = 0.168 mI4, Flow Rate = 0.3 L/min, 500 rpm.
243
-------
TABLE D-4. pH EFFECTS ON THE RATE CONSTANTS ku, and ko OF HCB
DESTRUCTION BY OZONATION AND PHOTOLYTIC OZONATION IN
PURIFIED WATER
PH
4
6
8
10
3.
3.
6.
1.
k *
Ko
56 x 1CT3
75 x 10~3
91 x 10~3
21 x 10~3
k **
Ku
0.89 x 10-!
1.40 x 10"1
2.12 x 10"1
* equation D-4
**equation D-5
244
-------
IOH
1 I • I
10 20 30 40
time , (min)
50
60
I
70
80
Figure D-l. Ozone concentration profile for ozonation of HCB
in lake water.
245
-------
2.0
1.5
D (mg/min-L)
O
A
D
1.14
2.16
3.04
3.95
o>
E
.0
O
**
z
O
0; 1.0
Z
LU
O
Z
O
O
UJ
z
O
N
O
0.5
20
30 40 50
TIME (min.)
60
70
80
Figure D-2.
Ozone concentration profile for photolytic ozona-
tion under 0.25 W/L UV radiation.
246
-------
1.0
[HCB]
[HCB].
.5
o
D
O
• © o
•5 A © ©
a
•
D A
0
A
Q
0
Q
Q
O
A
D
i 1
O
o
o
D (mg/min-L)
1.16
2.19
3.00
4.04
0 5 10
20
30
40
50
60
70
80
TIME (min)
Figure D-3.
Effects of dose rate on the direct ozonation of
HCB in lake water.
247
-------
T_ 3
1
DQ
00'
o
3.5
0
.0
1.5
InD ( mg/min- L )
Figure D-4.
Initial rates of direct ozonation of HCB in lake
water at various ozone dose rates.
248
-------
1.0
[HCB]
[HCBL
.5-H
D (mg/min-L)
0
A
1.14
2.16
D 3.04
• 3.95
JV = .25w/L
PREDICTED
O PHOTOLYSIS
ONLY
I"' I ' I ' .1 ' .1 ' J. ' 1 ' I ' 1
10 20 30 40 50
time , (min)
60
70
80
Figure D-5.
Effects of dose rate on the photolytic ozonation
of HCB in lake water.
249
-------
[HCB]
[HCBjo
0
.12
.25
.49
- PREDICTED
D= 3.04 mg/min-L
mi , i i
10 20
1 I ' I
30 40 50
time, (min)
60
70
I
80
Figure D-6.
Effect of UV radiation transfer rates on the photo-
lytic ozonation of HCB in lake water.
250
-------
1.0
[HCB]
[HCB]O
0.5 —
4.77
4.76
8 4.65
10 4.34
PREDICTED
1
10
I
20
I
30
I ' I
40 50
time, (min)
I
60
I
70
I
80
Figure D-7,
pH effects on the rate of ozonation of HCB in
purified water.
251
-------
[HCB]
[HCB]C
O 4 1.14
D 6 1.33
A 8 1.37
UV= .21 W/L
0.5-
10
20
30
40 50
time , (min)
60
70
BO
Figure D-8.
Effect of pH on rates of photolytic ozonation of
HCB in purified water
252
-------
[HCB]
[HCBJo
0.653
0.924
1.367
2.500
3.517
PREDICTED
pH = 8.0
UV= .21 W/L
0.5J
80
Figure D-9. Photolytic ozonation of HCB in purified water,
253
-------
1.0
HCB
0.5 -
O PURIFIED WATER
10
20
30
40
50
60
70
80
Figure D-10.
Comparison of photolytic ozonation in lake water
and purified water. UV intensity = 0.2 W/L.
2S4
-------
1.0
[HCB]
[HCB1
0.5
© PURIFIED WATER
A LAKE WATER
UV=.I w/L
D = .5 mg/min-|_
I
10 20 30 40 50
time , (mm)
60
70
80
Figure D-ll. Comparison of photolytic ozonation in lake water
and purified water. UV intensity = 0.1 W/L.
255
-------
APPENDIX E
IDENTIFICATION OF CHEMICAL BY-PRODUCTS FROM OZONE/
ULTRAVIOLET-INDUCED REACTIONS OF
ORGANIC COMPOUNDS IN WATER
EXPERIMENTAL PROCEDURES
I. Experimental Apparatus
Besides the large contacting reactor described in Section 5
above, a small reactor was designed particularly for the study
of oxidation by-products. Figure E-l shows the design of the
reactor. The body was made of quartz for the transfer of UV
radiation, and had a reaction volume of 45 mL. Ozone was
directed through the Pyrex tubing and sparged into the solution
by a glass frit. The whole reactor was placed at the center of
the Rayonet Photochemical Reactor. The operation of the reaction
system was similar to that with the large contacting reactor,
but the flow of 03/02 was kept at 75 ml/min. The Swagelok nuts
on the ozone inlet could be removed for taking samples from the
solution; however, in so doing, the supplies of 03 and UV were
temporarily interrupted.
Analytical Instruments
Gas Chromatography/Mass Spectrometry
The Finnigan Model 3200 gas chromatograph/mass spectrometer
(GC/MS) system with a Model 6100 data system was used to separate,
detect and identify the components. All samples were chromato-
graphed using a 30 meter by 0.25 mm I.D. SE-30 wall coated
glass capillary column (Supelco, Inc.). A splitless technique
was used in order to obtain high detectability. All the data
were stored on magnetic tapes and interpreted by comparison with
published mass spectra, commercially available computer data or
by a priori interpretation based on the mass fragments. The
GC/MS system was used primarily for the identification of the
structure of products. Some packed columns were also used tp
assist in the identification of peaks found using the Model
5710A GC/FID.
256
-------
Figure E-l. Reactor for ozone and ozone/UV treatment of HCB.
257
-------
High Pressure Liquid Chromatography (HPLC)
The Waters Associates Model ALC-201 liquid chromatograph
and solvent programmable Micromeritics Model 7000B liquid
chromatograph were used to separate large amounts of product
samples and to perform fractional collection for GC/MS identi-
fication. A Tracor 974A variable wavelength detector was
connected to either liquid chromatograph. A wavelength of 255 nm
was used for all separations. Organic solvents used were
filtered and degassed. Purified water described previously
was pumped through a C-18 reverse phase column for further
purification. Three types of HPLC columns were used, one
containing Partisil 10 ODS (permanently bonded C-18) (Whatman),
another containing Bondapak C-18/Corasil, 35-50 microns (Waters
Associate, Inc.), and the third containing Lichrosorb RP-18, 10
microns (EM Laboratories, Inc.). All were slurry'packed into
15 cm by 4.6 mm I.D. stainless steel columns in the laboratory.
An electrochemical detector (BCD), Model LC-4 (Bioanalytical
Systems Inc., West Lafayette, Ind.), was also used for detection
of easily oxidized or reduced components in freeze-drying
samples from the reaction mixture. In most cases, BCD and UV
detectors were used and compared simultaneously.
Microcoulometry
A Dohrmann Envirotech Microcoulometric Titration System
(MCTS-20) was used to measure the amount of chloride in the
reaction mixture. The sample was injected directly into the
T-300-S titration cell. The amount of chloride in nanograms
was shown digitally on the display on the front panel.
Analytical Procedures
HCB in methanol was spiked into 45 mL of purified water as
a suspension; usually 2-5 mg was applied. It was found
necessary to make product runs at a much higher concentration
than that used in the kinetic runs and, since the solubility
limit of these relatively insoluble compounds was greatly
exceeded, a cloudy suspension of substrate was the starting
material for these studies. After ozone/UV treatment for a
brief period, the cloudiness disappeared indicating that the
substrate had reacted, and the treatment was stopped. The
suspension was subjected to photolysis, ozonation and photolytic
ozonation. Ozone dose rate was 6.14 mg/min. At the end of a
run, the whole solution was extracted with anhydrous ether
according to the scheme shown in Figure E-2. The acidic
extraction fraction was methylated using diazomethane by the
procedure and apparatus of Keith (158). Both neutral and
acidic fractions were concentrated to a final volume of 200 yL
and internal standard, hexamethyl benzene, was added for
quantitative determination of the yields of products.
258
-------
SAMPLE
E(2°
DRY w/ Na2S04
ir
KUDERNA DANISH
CONCENTRATION
ADJUST
pH=ll
i
ADJUST
pH = 3
EtgO
1
DRY w/ No2S04
KUDERNA DANISH
CONCENTRATION
METHYLATION
with
CH2N2
i
GC/MS
Figure E-2. Extraction scheme.
259
-------
The products were chromatographed on the Hewlett Packard Model
5710A GC/FID or Finnigan Model 3200 GC/MS. Structures were
assigned by rationalization of the mass spectra. These were
confirmed by comparison with known spectra from the literature
(164-166) or from analysis of an authentic sample in this
laboratory.
Nonextracted small acids and oxalic acid were determined
by the modified methods of Bethge and Lindstrom (159). A portion
of aqueous solution resulting from photolytic ozonation was
made basic by adding tetrabutylammonium hydroxide and dried at
40°C in a heated vacuum desiccator, Precision Scientific Co.,
Chicago, Illinois. The residues were dissolved in acetone and
either benzyl bromide or an alkyl iodide was added. Small acids
were converted to benzyl derivatives and GC/FID chromatographed
on the butanediol succinate column. Oxalic acid was converted
to its ethyl derivative and chromatographed on the SP-1000/H3P04
column.
The relationship of the formation of several major products
to the photolytic ozonation time were studied. Several runs
were carried out, each having different reaction time, and the
products in each sample were chromatographed by GC-FID and
GC/MS using the SP-2100 column and SE-30 capillary column.
The large contacting reactor was also used for study of
products. Three liters of reaction solution were extracted
according to the scheme in Figure E-2. The products were
identified by GC/MS and the concentrated samples were chromato-
graphed using HPLC. Collections of HPLC fractions was carried
out by collecting the components in 5 mL vials immediately after
the UV detection cell.
In the case of the volatile chlorinated compounds chloroform.
bromodichloromethane, and tetrachloroethylene, the chromato-
grams obtained from the kinetic work were searched for the
appearance of new peaks during a run, which would probably
indicate that a halogen was being formed. In only one case was
such a peak found during ozone/UV or ozone-only experiments,
and those results will be described in the later section which
covers results.
In addition to chromatographic monitoring of the reaction
mixture, product runs were made for chloroform and tetrachloro-
ethylene in which the off-gas was bubbled through basic
solution in order to trap expelled carbon dioxide. The off-gas
was next directed to another trap containing pentane or
isooctane, in order to trap volatile nonpolar compounds. When
gas chromatographic analysis showed that most of the substrate
had been destroyed in a particular run, additional gas was
sparged through the headspace to ensure the removal of C02 to
the trap, then the reaction mixture was analyzed by
260
-------
microcoulometry to determine the chloride content and by
titration to determine the amount of hydrogen ion formed.
Ozone consumption was calculated from the dose rate and
treatment time. The final reaction mixture was extracted by a
modification of the method of Henderson, Peyton and Glaze (167)
and that extract checked for the presence of products by gas
chromatography and GC/MS.
Ozonation and Photolytic Ozonation Products of Natural Water
Caddo Lake water was ozonated without UV irradiation in
order to identify products of the oxidation/photolysis reactions
which occur. Three liters of filtered Caddo Lake water was
treated for 25 minutes with 500 ml/min of 2.4% by weight of
ozone in oxygen, corresponding to a dose rate of 28.4 mg/min.
The reactor was stirred at a rate of 365 rpm, and in the UV run,
8 low-pressure, 75 W bulbs were used. One liter of the treated
water was adjusted to pH 11 and extracted with 150 ml, 75 ml,
and 50 ml portions of ether. The extracts were passed through a
drying column of sodium sulfate and then concentrated to 200 yl.
The extracted water was adjusted to pH 3, extracted and dried as
above, and concentrated to 2 ml. This was split into two
portions, one being derivatized with Methelute and one with
diazomethane. Controls (no ozone) and ether and diazomethane
blanks were run, and these, as well as the samples, were
diluted with an equal volume of hexane prior to splitless
capillary chromatography/mass spectrometry using a 30 meter
SP-2100 WCOT column in a Finnigan 3200 GC/MS system. The
samples derivitized with methelute had many spurious and inter-
fering peaks which were present in the control, and those
samples were discarded.
I. Results of Product Studies on Chloroform
These results are given in the body of the report.
II. Results of Product Studies on 2,2',4,4',6,6'-Hexachloro-
biphenyl
This section of Appendix E contains reconstructed gas
chromatograms and mass spectra from which products of the ozone/
UV treatment of hexachlorobiphenyl were identified. Figures E-3
-E-7 show the reconstructed gas chromatograms and mass spectra
from product studies on HCB photolysis. Table E-l contains the
major acidic products identified and their location in Figures
E-8, E-9, and E-10. Figures E-ll through E-17 give mass
spectra and major fragmentation patterns for products identified
in the reconstructed gas chromatogram shown in Figure E-8.
Results are discussed in Section 6 of the main report.
261
-------
100
cn
to
ISO
50
1.00
- | - 1 - !
150
200
250
450
0 550 600 650 700 750 800 850 900
Figure E-3. GC/MS scan of HCB after photolysis under nitrogen.
-------
to
IIII I 11 I 11 III 111 111 I III I 1111 I I 111 I 1111111III 11 11 111III I I I I I I III l-l
650 TOO TbO 800 850 900 950
Figure E-4. GC/MS scan of HCB after photolysis under oxygen.
-------
100
ro
50
100
300
100
150
200
250
3i50 400 450
Figure E-5. Mass spectrum of spectrum no. 834.
-------
100
fO
J IJ ll
50
100
100
(iii ill I
150
200
250
300 350 400
Figure E-6. Mass spectrum of spectrum no. 806.
-------
100
100
150
200
250
to
100
' 3bO
3l50'
Figure E-7. Mass spectrum of spectrum no. 903.
-------
HCB/03/UV RCIDIC FRflCTION CONC. 70 EV
100
0 0
CH30-C-CH2-CH2-C-OCHj
120 150 200 250 300 350 400 450 500 550 'f'600
100
MTI > M 1 ' I 1 1 t I i'l i'1 * I I'TI 1 I I I I I I i IJ I ' I i 1 ' I i M I
TOO 750 800 850 900
620
(A)
0
C-OCH3
0»c-OCH3
(F)
(K)
0
C-OCH3
(B)
(0
(D)
C-OCH3
0
C-H
0
C-OCH3
Ifi3
°*C-OCH3
(E)
(G)
°»C-OCH3
*0
CHs0.c
0VMH8
(I)
(J)
00VOCH3
CH30-C
[O
CHjO
CH30-C>
(L)
(M)
OH 0
(O
Q.C-OCH3
0
C-OCH3
0
C-OCH3
(N) Y$)
Figure E-8. GC/MS scan of acid fraction from the photolytic
ozonation of HCB.
267
-------
o OYOCHJ
'i
00
Figure E-9. FID chromatogram of HCB/03/UV reaction mixture,
-------
Figure E-10,
I 2
Reaction Time. -hours
HCB photolytic ozonation products at various time
intervals.
269
-------
100
:Ui
60
,li
i
100
150
200
250
144
-Cl
Figure E-ll. Mass spectrum of spectrum no. 590.
-------
100
NJ
100
150
200
250
OCH.
10
m/e = 252
-Cl
217
Figure E-12. Mass spectrum of spectrum no. 640.
-------
100
JKU-r,
> 250 X 10
10
100
300
100
150
200
250
10
m/e= 280
172 T245
Figure E-13. Mass spectrum of spectrum no. 690.
-------
100
100
150
200
350
> 250 X
250
-CH2
-HCl
-28
m/e= 298
'209 "IV" 7173
Figure E-14. Mass spectrum of spectrum no. 700.
-HCl
109
-------
100
15
to
100
277
Figure E-15. Mass spectrum of spectrum no. 755.
-------
ro
100
150
200
250
C
ii
C
168
-Cl
400
297
Figure E-16. Mass spectrum of spectrum no. 807.
-------
-59
IO
-CH ^SX^Q -CH3OC-OH
r*t> CH30
0° o'
0
CH,0
325 "CH2 T339
335
Figure E-17. Mass spectrum of spectrum no. 820.
-------
TABLE E-l. MAJOR ACIDIC PRODUCTS IDENTIFIED IN PHOTOLYTIC
OZONATION OF HCB IN WATER
Peak No.#
Compound ( )*
Structure
1.
2.
3.
2,4,6-trichlorobenzoic acid
(590)
2- (2,4,6-trichlorophenyl)
acetic acid (640)
3-keto-3- (2,4,6-trichlorophenyl)-
propanoic acid (690)
4.
2-chloro-3-(2,4,6-trichlorophenyl)-
2-propenoic acid (700)
# Peak number from Figure E-9
( )* Spectrum number of compound from Figure E-8
277
-------
Peak No.t
TABLE E-l (Continued)
Compound ( )*
Structure
5.
2- (2,4-dichlorophenyl)-3-
chloro-2-pentenedioic acid
(755)
2-chloro-3-(2,4,6-trichloro-
phenyl) maleic acid (804)
2-(2,4,6-trichlorophenyl)-3-
chloro-2-pentenedioic acid
(820)
OH
* Peak number from Figure E-9
( )* Spectrum number of compound from Figure E-8
278
-------
III. Results of Product Studies on Caddo Lake Water
Following, in Figures E-18 through E-23 and Table E-2 are
the reconstructed gas chromatograms and corresponding list of
compounds determined in extracts from controls, ozonated Caddo
Lake water and the same water treated with ozone/UV. Results
are discussed in Section 6 of the main report.
27'9
-------
100
A
N)
CO
O
100
A
'I"1""11!"
50 100
J » 1 I I 1 1 1 I 1 J 1 T ' I ' I ' 1 ' J ' 1 • 1 ' I ' 1 ' ] ' 1 • 1 ' I ' 1 ' ] ' 1 ' 1 T I ' I ' ] ' 1 '
150 200
250
300
350
400
:
AJi
"I"" I
450
sbo BOG ebo TOO
'dso sbo gfeo
Figure E-18. Reconstructed gas chromatogram; Caddo Lake water, ozonated. netural
fraction.
-------
to
CO
100
50
100
450
330
i'A
!M!>
550
600 61
u
50 700
n
1 i
750 800
p
|.
850
p
I
900
p
950
Figure E-19. Reconstructed gas chromatogram; Caddo Lake water, ozone/UV. neutral
fraction.
-------
100
00
to
100
""I1"
250
rT
300
350
400
450
A
doo""" 550 ebo eso
750 800 850 900
9*50
Figure E-20. Reconstructed,gas chromatogram; Caddo Lake water, control. neutral
fraction.
-------
100
to
S
200
250
300
350
400
1 I
m 1
450
650 700 750
850 900
Figure E-21. Reconstructed gas chromatogram; Caddo Lake water, ozonated. acidic
fraction. (diazomethane)
-------
100
A_
"i^*^11"1 i"
50 100
150
200
250
300
350
400
450
NO
00
100
w i w
It
b1
500 550 600
660
750 ebo
9^6
Figure E-22.
Reconstructed gas chroma tograra; Caddo Lake water, ozone/UV. acidic
fraction (diazomethane) m
-------
100
50
100
150
200
250
300
350
400
450
10
00
Ul
100
550 000 650 700 750 800 850
Figure E-23. Reconstructed gas chromatogram; Caddo Lake water, control.
fraction (diazomethane).
960
acidic
-------
TABLE E-2. PRODUCTS IDENTIFIED FROM THE OZONE AND OZONE/UV
TREATMENT OF WATER FROM CADDO LAKE, TEXAS
Method of
Identifi- Occurs in
Product cation Chromatogram
a
b
c
d
e
f
g
h
i
j
k
1
m
n >
P
q
r
s
t
V
w
X
y
z
n-Heptane
2 , 4-Dimethylhexane
Note 4a, below
Note 4b, below
Note 4c, below
Trichlorobenzene ,
internal standard
2 , 3-Dimethylheptane
n-Decane
Dicyclohexyl Amine
2 , 6-Di-t-Butylbenzoquinone
Note 4d, below
M=195, similar to substituted
pyridones
Diethyl Phthalate
§Hno-<0>~~<0^CH'£oXg>
Hydrocarbon
Bleed
Phthalate
3-Methylheptane
2 . 6-Di-t-Butyl-4-Ethyl
M.e Phenol
N^N
N ]
s\
Me
N,N-Diethylformamide
Note 4e, below
Propyl Ethyl Ketone
Methyl Butyl Ketone
see Note 4f, below
1
1
2
-
-
1
1
1
1
1
-
2
1
3
2
2
2
1
1
1
1
-
1
1
-
I, II, III, IV (Note 5)
I
I, II, III
I, II, III
I, II, III
I, II, III
I
I
I, III, IV
I
I
I
I, III
I, II, III
I, II, III
I
I
II, III, IV
I
II
II
II, III, IV
III
III
III
(Continued)
286
-------
TABLE E-2 (Continued).
a1
b1
c1
d1
Product
RC02Me
RCOJYle, R >C,
2. b
CH3(CH2)12C02Me
CH3(CH2)14C02Me
Method of
Identification
2
2
1
1
Occurs in
Chroma togram
III
III
III
III
1)
Registry of Mass
spectral Data, Stenhagen, E.;
Abrahammson,
S.; McLafferty, F.W.: John Wiley, New York, 1974.
2) Suggested by fragmentation pattern.
3) Fetizon, M. and Audier, H. : Organic Mass Spectrometry, 8_
(1974) 201.
4) No positive identification. Distinctive ions along with
suggested compounds and source are given below:
CH,OCH-CH=CH,*«-^CH,OCH-CH-CH,
a i 2 a \ / o
OH 0
.
Acrolein Methyl Hemiacetal M/e=88 , 73 , 70 , 61, 45 , 44^, 42
b)
c)
CH8
CH3CH2CN=N-OH or
M/e=85> 84, 57, 43
,CH^CH,NO M/e=73, 55, 45
d) Almost matches Registry spectrum for Delta-10-
Tetrahydroconstunolide,
but not quite. Difference
could be the instrument, but spectrum is similar in
fragmentation pattern to
287
-------
OCH,
+16.
Possibly
or
5)
0 OH
The latter are often used as antioxidants in ether.
e) Long chain N-Methyl Amine or possibly Hydrocarbon.
M/e = 85 < 84, 57, 55, 43
f) M/e = 111, 89, 75 < 73, 47 < 45 > 43
I = ozone only: neutral and basic compounds; II = ozone/UV:
neutral and basic; III = ozone only: acidic compounds
(Diazomethane derivatized); IV = ozone/UV: acidic compounds
(Diazomethane derivatized).
288
-------
APPENDIX F
QUALITY CONTROL DURING THE KINETIC RUNS
The purpose of this appendix is to outline sources of
experimental error which were unique to this project, present
experimentally determined limits to these errors, and evaluate
the importance of these errors on the desired outcome. There
are several difficulties which combine to produce the unique
experimental environment in which an analytical method must be
chosen for this project. The first is the low concentration
level at which these experiments were run in order to realis-
tically simulate an actual water supply. This is, of course,
more of a problem in some cases than others but more importantly,
combines synergetically with the following. The ozone/UV system
is highly complex, highly reactive, and not well understood. The
effects the complexity can have on the analytical method are
unknown, and after consideration of all known effects, one can
only hope that the data are real, as is the case with any
experimental endeavor. The simpler and faster the analytical
method, however, the less opportunity there is for artifacts to
occur. The greatest problem in this particular project is the
quantity of data which must be collected to achieve the desired
result. Fifteen to twenty gas-chromatographic analyses may go
into determining one disappearance curve. Multiplying by the
number of kinetic runs performed during this project gives a
very large number of analyses. Doubling or tripling the number
of analyses by replication doubles or triples the length of the
project without significantly increasing the amount of infor-
mation which is gained. The result of this trade-off is that
it is desirable to use analyses which only require one injection.
This choice also satisfies the simplicity and speed requirements
mentioned earlier. Next to direct monitoring (such as spectro-
scopic) inside the reactor, the natural choice for the analytical
method would be direct aqueous injection gas chromatography,
where allowed by detection limits, since sample handling is
minimized and the analysis is performed immediately upon
sampling (within 30 seconds). Analyses were performed in this
manner for chloroform, bromodichloromethane, and tetrachloro-
ethylene, but not for hexachlorobiphenyl. Hexachlorobiphenyl
requires a concentration step and is therefore extracted from
ozone-quenched aliquot of reaction mixture, using hexane to which
an internal standard has been previously added. Data for
specific substrates will now be considered by individual compound
289
-------
The concept of standard deviation (N-l weighted) will be
employed as a measure of precision, even though a statistically
valid number of data points has not been collected.
CHLOROFORM
Standard solutions spanning the concentration range
0-400 yg/lTwere made up by spiking known concentrations of an
ethanol concentrate into water. The most concentrated solution
was used to calibrate the computerized Hewlett-Packard 5840 gas
chromatograph, then the remaining solutions were analyzed on the
basis of that calibration. Only one injection was made per
solution to simulate the analysis of an actual kinetic run. The
results are shown in Figure F-l and in the following table.
Prepared concentration,yg/L 160 80.0 40.0 20.0 4
Analyzed concentration,yg/L 162 73.8 36.8 19.5 n.d.
The purified water was checked for chloroform by the same method,
and it was found to be negligible. The peak corresponding to
4 yg/L was visible but was not integrated by the electronic
integrator. Another linearity study was performed similarly
over a narrower range, and duplicate injections were made. The
integrator was calibrated on the highest concentration. The
data is shown in Figure F-2 and this table:
Prepared concentration,yg/L 37.1 23.2 18.5 1.3.9 9.3
Analyzed concentration^ g/L 37.1 24.6+. 05 17+2 13.2+.9 9.5+5
In another experiment, three separate solutions of aqueous
chloroform were prepared from the same concentrate, to determine
the variance of Co (substrate concentration at t=0) from one
run to the next. Each solution was analyzed in triplicate and
the averages averaged in order to calculate a standard deviation
among the solutions. The data is shown below:
Peak Areas, yV Seconds -f- 3. 2^ Average "LStd. Dev.
Solution A 162,300 163,500 161,400 162,400+1053(0.6%)
Solution B 154,800 154,500 152,700 154,000+1136(0.7%)
Solution C 153,100 155,600 157,100 155,266+2021(1.3%)
Average of Averages 157,222+4528(2.9%)
f The Hewlett-Packard integrator uses units of 3.2 yV Seconds
Solution B was then diluted by a factor of ten and that
solution analyzed in triplicate:
1/10B 14,210 15,950 12,700 Avg=14,286 + 1626 (11.4%)
290
-------
These figures indicate that at either Co = 100 yg/L or
1/10 Co = 10 yg/L, the precision of the measurement is about
1 yg/L.
BROMODICHLOROMETHANE
Solutions of five different concentrations in the range of
3-300 yg/L were prepared by spiking an ethanol concentrate into
water, and each resulting solution analyzed by a single direct
aqueous injection. This time the integrator was calibrated on
the lowest concentration:
Prepared concentration, yg/L 2.93 14.6 29.3 147 293
Analyzed concentration, yg/L 2.93 12.4 28.9 149 286
The data are plotted in Figure F-3. As with chloroform,
three aqueous solutions were made up identically to be 100 yg/L
in bromodichloromethane. A portion of one of the solutions was
diluted by a factor of ten, and a portion of the resulting
solution again diluted by a factor of ten. The resulting
solutions were analyzed in triplicate. The data are given below
and plotted in Figure 4.
Peak Area, yV Seconds f 3.2 Average ± Std. Dev.
Solution A 719,800 716,600 724,000 720,133+3711 (0.5%)
Solution B 686,200 693,400 695,800 691,800+4995 (0.7%)
Solution C 675,600 662,400 675,400 671,133+7564 (1.1%)
Average of Averages-694,355+24,599 (3.5%)
1/10 B 54,570 55,040 55,020 54,877+266 (0.5%)
1/100 B 4,717 4,758 4,680 4,718+39 (0.8%)
Comparison of the precision of analysis to the reproduci-
bility of solution preparation for solutions A, B, and C, seems
to indicate that the standards can not be prepared as accurately
as they could be analyzed. This factor is probably the major
contributor to the deviation on the 10 yg/L point from the
line in Figure F-4. This deviation corresponds to about 2 yg/L.
HEXACHLOROBIPHENYL
Three solutions of HCB were identically prepared by spiking
from an ethanolic concentrate. The solutions which were made up
to contain 60 yg/L of HCB, were analyzed in triplicate by
extraction into nanograde hexanes (mixture of isomers) to which
had been added hexachlorobenzene as an internal standard.
Because of the instrumental work load, a gas chromatograph
which did not have a functioning integrator was used for the work.
291
-------
The integrator available was able, however, to measure peak
heights electronically, and the results reported here are based
on those heights. One of the solutions was diluted by a factor
of six and a portion of that solution diluted by another factor
of six. These two solutions were also analyzed in triplicate as
above. The data are given in the following table and are shown
in Figure F-5.
Solution
A(60 yg/L)
B
C
Peak Height HCB/Peak Height Internal Standard
#1
#2
#3
1.312 1.483 1.498
1.507 1.557 1.378
1.526 1.436 1.435
Average ± Std. Dev.
1.43 + 0.10 (7.0%)
1.48 + 0.09 (6.1%)
1.47 + 0.05 (3.4%)
Average of Averages 1.46 + .03 (1.8%)
1/6 A(10yg/L) 0.256 0.246 0-237
1/36 A(1.65yg/L) 0.0571 0.0559 0.0563
0.249 + 0.009 (3.6%)
0.0564+0.0006 (1.1%)
In contrast to the case of bromodichloromethane, these data
seem to indicate that the solutions (Co) can be prepared with
reasonable accuracy (+1.8%), but that the analytical method is
not as precise as for the two compounds described previously.
This is to be expected since there are three measurement steps
(sample, quench, and extraction solvent), and one extraction,
rather than just one measurement as in direct aqueous injection.
Measurement of the injection volume into the syringe is not
counted for the HCB analysis since this is an internal standard
technique.
EFFECT OF ANALYTICAL PRECISION ON END RESULTS
At the present time only the effect on T will be discussed.
It will, further, be discussed only for the case of a reaction
which is first order in substrate, in the interest of simplicity.
When integrated, the rate expression gives
In S/SQ = -kt.
Let -rU) be the time at which S/SO = e~A.
If the standard deviation is expressed as a fraction, e,
of the measured value, then the limits set by the above equation
on t correspond to the case in which the errors are in opposite
directions.
t + = - =•
d+e). =
= -i [In 2_
In
(1 + e)
TT-r^>
or
292
-------
6 + = t - t + = 1 . 1+
~ ~
o
The fraction by which this measurement should deviate from the
true value of T is then
(1) = 6+ /T (1) = lnj-§,/ In (S/So), but
+
E^
+£
0 +
In S/So = In e'1 = -1, or #+(l) = 6+ /T (1) = =- In f-E
L O
For example, the analysis of solution A of HCB gave e =0.07.
Using the percent deviation of solution 1 A (17% remaining) as
an approximation to T(3) (5% remaining) , then e = 0.036, and
= - - In i = -0.108, J-(i) = 0.104
/+(3) = 0.036 , #-(3) = 0.035.
In other words, the lack of precision encountered in the
analysis of solution A of HCB could contribute about 10% error
to T(l) and about 3^% to T(3). If, on the other hand, the
error from 1/36 A had been used, the effect on T would be
f+(l) = 0.084 #-(1) = 0.079
= 0.028 #-(3) = 0.026.
It can be seen from the sample calculations given above and
below that ^±d)2ie+eo and ^(!) = -r ^(D • This is easily proven
using the series expansion of In (1 + x) .
Four important points should be made about the above effect
1) #+(1) represents the worst possible case, that is,
maximum errors in opposite directions.
2) Some of the worst data obtained in the precision
studies was used in the above tables; calculations
with e = e0 = 0.036 yield f +(1) = 0.072 and
#+(3) = 0.024.
3) The importance of the effect is less as more substrate
is destroyed.
293
-------
4) The fact that the standard deviations for the indi-
vidual solutions in the HCB experiment are much higher
than that for the average of the averages, implies
that if one fits a destruction curve to the data,
and then extracts T from that curve, the actual error
in T will be smaller than that predicted by the above
analysis.
As can be seen in Appendix B, the scatter in the
kinetically determined rate constants has been calculated as
a standard deviation and reported in the summary tables in
Section 5.
294
-------
400r
100 200 300
CHLOROFORM, PREPARED, pg/l
400
Figure F-l.
Linearity of chloroform analysis,
(0-400 yg/L).
Broad Range
295
-------
10 20 30
CHLOROFORM , PREPARED,
40
Figure F-2.
Linearity of chloroform analysis,
(o-40 yg/L).
Narrow Range
29.6
-------
300r
100
BROMODICHLOROMETHANE
200
PREPARED,
300
Figure F-3. Linearity of bromodichloroniethane analysis,
297
-------
8.0
6.0
40
o
x
CO
<
uj
o:
2 2.0
CL
20 40 60 80
BROMODICHLOROMETHANE , pg/l , AS PREPARED
100
Figure F-4 .
Linearity of bromodichloromethane analysis upon
serial dilution in water.
298
-------
1.5 r
Q
tr
to 1.0
o:
UJ
m
o
X
Q 0.5
fe
o:
UJ
X
UJ
Q.
10 20 30 40
HCB,|jg/l,AS PREPARED
50
60
Figure F-5.
Linearity of hexachlorobiphenyl analysis upon serial
dilution in water. (error bars shown for solution
A only).
299
-------
TECHNICAL REPORT DATA
(Please read Instructions on the reverse before completing)
1. REPORT NO.
EPA-600/2-80-110
3. RECIPIENT'S ACCESSION-NO.
4. TITLE AND SUBTITLE
OXIDATION OF WATER SUPPLY REFRACTORY SPECIES
BY OZONE WITH ULTRAVIOLET RADIATION
5. REPORT DATE
August 1980 (Issuing Date)
6. PERFORMING ORGANIZATION CODE
7. AUTHOR(S)
William H. Glaze; Gary R. Peyton; Francis Y.
Huang; Jimmie L. Burleson; Priscilla C. Jones
'8. PERFORMING ORGANIZATION REPORT NO,
9. PERFORMING ORGANIZATION NAME AND ADDRESS
Institute of Applied Sciences
North Texas State University
N. T. Box 13078
Denton, Texas 76203
10. PROGRAM ELEMENT NO.
61ClCfC824A,SOSl,Task32
11. CONTRACT/GRANT NO.
R-804640
12. SPONSORING AGENCY NAME AND ADDRESS Cin. , OH
Municipal Environmental Research Laboratory-
Office of Research and Development
U.S. Environmental Protection Agency
Cincinnati, Ohio 45268
13. TYPE OF REPORT AND PERIOD COVERED
Final. 9/76 to 2/80
14. SPONSORING AGENCY CODE
EPA/600/14
15. SUPPLEMENTARY NOTES
Project Officer: J. Keith Carswell (513) 684-7228
16. ABSTRACT
The use of ozone with ultraviolet radiation was studied as an advanced
treatment process for the removal of micropollutants and trihalomethane
precursors from drinking water. The model compounds chloroform, bromo-
dichloromethane, tetrachloroethylene and 2,2",4,4",6,6"-hexachlorobi-
phenyl were treated with ozone/UV as well as UV and ozone individually
in both highly purified water and lake water. Kinetic rate expressions
which express the dependence of the reaction rate on ozone dose rate, UV
intensity and substrate concentration were developed. The trihalomethan<
formation potential of a natural lake water was monitored as a function
of ozone/UV treatment time. Products resulting from the ozone/UV treat-
ment of some model compounds and lake water were studied. The kinetic
data were submitted to a subcontracted consulting engineering firm for
full-scale process design and a treatment cost estimate. Costs for a
1-50 Mgal/day plant to remove 90% of 100 ug/L of chloroform and bromodi-
chloromethane and 20 ug/L of tetrachloroethylene were $0.063 - 0.16 per
thousand gallons. Costs for removal of 90% of the trihalomethane for-
mation potential from lake water were $0.099 - 0.20 per thousand gallons
17.
KEY WORDS AND DOCUMENT ANALYSIS
DESCRIPTORS
b.lDENTIFIERS/OPEN ENDED TERMS
c. COSATI Field/Group
Oxidation
Ozone
Ultraviolet radiation
Water treatment
Cost estimates
Precursors
Trihalomethane
Chemical kinetics
Design scale-up
13B
8. DISTRIBUTION STATEMENT
Release to Public
19. SECURITY CLASS (ThisReport)
Unclassified
21. NO. OF PAGES
320
20. SECURITY CLASS (Thispage)
Unclassified
22. PRICE
EPA Form 2220-1 (9-73)
300
ffU.S GOVERNMENT PRINTING OFFICE: 1980-657-165/0113
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