• TABLE OF CONTENTS
HIDE
 Title Page
 Acknowledgement
 Table of Contents
 Abstract
 Introduction
 Materials and methods
 Degradation of Alkyl Halides
 Solubilization and degradation...
 Gasoline in ground water
 Summary and conclusions
 Solubility measurements by Linda...
 Fortran program for modeling loss...
 Area count data set for statistical...
 Reference
 Biographical sketch
 Copyright














Title: Behavior of partially miscible organic compounds in simulated ground water systems
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Permanent Link: http://ufdc.ufl.edu/UF00086020/00001
 Material Information
Title: Behavior of partially miscible organic compounds in simulated ground water systems
Physical Description: vii, 194 leaves : ill. ; 28 cm.
Language: English
Creator: Cline, Patricia V
Publication Date: 1988
 Subjects
Subject: Groundwater -- Pollution   ( lcsh )
Trichloroethane -- Solubility   ( lcsh )
Gasoline -- Solubility   ( lcsh )
Environmental Engineering Sciences thesis Ph.D
Dissertations, Academic -- Environmental Engineering Sciences -- UF
Genre: bibliography   ( marcgt )
non-fiction   ( marcgt )
 Notes
Thesis: Thesis (Ph. D.)--University of Florida, 1988.
Bibliography: Includes bibliographical references.
Statement of Responsibility: by Patricia V. Cline.
General Note: Typescript.
General Note: Vita.
 Record Information
Bibliographic ID: UF00086020
Volume ID: VID00001
Source Institution: University of Florida
Rights Management: All rights reserved by the source institution and holding location.
Resource Identifier: aleph - 001116907
oclc - 19924511
notis - AFL3692

Table of Contents
    Title Page
        Page i
    Acknowledgement
        Page ii
    Table of Contents
        Page iii
        Page iv
    Abstract
        Page v
        Page vi
        Page vii
    Introduction
        Page 1
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    Materials and methods
        Page 12
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    Degradation of Alkyl Halides
        Page 22
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    Solubilization and degradation of residual TCA
        Page 75
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    Gasoline in ground water
        Page 95
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    Summary and conclusions
        Page 166
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    Solubility measurements by Linda Lee
        Page 171
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    Fortran program for modeling loss of residual TCA
        Page 173
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    Area count data set for statistical analysis of water extracts of gasoline
        Page 179
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    Reference
        Page 187
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    Biographical sketch
        Page 194
        Page 195
        Page 196
    Copyright
        Copyright
Full Text












BEHAVIOR OF PARTIALLY MISCIBLE ORGANIC COMPOUNDS
IN SIMULATED GROUND WATER SYSTEMS

















BY

PATRICIA V. CLINE


A DISSERTATION PRESENTED TO THE GRADUATE SCHOOL
OF THE UNIVERSITY OF FLORIDA IN
PARTIAL FULFILLMENT OF THE REQUIREMENTS
FOR THE DEGREE OF DOCTOR OF PHILOSOPHY

UNIVERSITY OF FLORIDA

1988


i l Sy I OF FORIDA LIBRARIES















ACKNOWLEDGMENTS

I sincerely appreciate the technical and editorial

assistance provided by my research director, Dr. J. Delfino.

I also thank Dr. P. S. C. Rao and Dr. P. Chadik for their

advice and for providing opportunities for challenging

discussions, and Dr. J. Dorsey and Dr. R. Yost for serving

on my committee and reviewing this dissertation. Each

member of my committee has contributed to my graduate career

through excellent teaching and creating a positive

intellectual environment.

This work was funded by grants from the Florida

Department of Environmental Regulations. Special thanks are

extended to Dr. Geoffrey Watts for his role in securing

funds and providing technical support and comments.

I am grateful to Dr. M. Battiste for discussions of

reaction mechanisms, and for providing the use of his

laboratory for the synthesis of brominated ethanes.

Special thanks go to Angle Harder for her hard work,

Linda Lee for her generosity with analyses and information,

Tom Potter for unselfish computer and mathematical

assistance, and Bill Davis for technical support.

I extend warmest and deepest thanks to my husband Ken

for technical assistance and emotional support, and my son

Brendan for giving me joy.















TABLE OF CONTENTS


ACKNOWLEDGEMENTS .


ABSTRACT .

INTRODUCTION


Chemistry of Alkyl Halides
Gasoline Partitioning .


MATERIALS AND METHODS .. ..


Alkyl Halides
Gasoline .


DEGRADATION OF ALKYL HALIDES


Introduction . . . . .
Degradation of 1,1,1-Trichloroethane . .
Degradation of Brominated Ethanes . .
Halogenated Ethenes . . . .
Structure/Rate Relationships of Alkyl Halides .
Simple SN1/E1 Reactions . . .
Comparisons of Geminal Trihalides . .
Effect of Additional Halogens on the Alpha
Carbon . . . . .
Sediment Matrix Affects . . . .


SOLUBILIZATION AND DEGRADATION OF RESIDUAL TCA .

Behavior of Residual Solvent . .
Aqueous Phase Concentrations . .
Advection . . . . .
Degradation Rate . . . .
Model Parameters and Procedures . .
Limitations of the Model Assumptions ..


iii


. . . ii


. . . 1


S. 3
. 8


. . . . 1 2
. . . . 18


. . 22











GASOLINE IN GROUND WATER


Background . . . . 95
Composition of Gasoline . .. 95
Multicomponent Liquid-Liquid Equilibria 97
Statistics and Pattern Recognition
Applications . . . 102
Partitioning of Gasoline Components into Water 105
Fuel/Water Partition Coefficients . .. 105
Water Soluble Blending Agents .. ... 113
Prediction of Kfw for Other Components 120
Changes in Concentrations with Time .. 122
Differences in Water Extracts of Gasolines 129
Equilibrium Concentrations of Major
Constituents . . . .. 131
Visual Comparison of Water Extracts of
Gasoline . . . .. 131
Preparation of the Data Base for Statistical
Analysis . . . . 135
Basic Descriptive Statistics . .. 138
Bivariate Plots . . . 141
Stepwise Discriminant Analysis . .. 148
Principal Component Analysis . .. 155
Summary . . . . 164

SUMMARY AND CONCLUSIONS . . . 166

APPENDIX A. SOLUBILITY MEASUREMENTS BY LINDA LEE 171

APPENDIX B. FORTRAN PROGRAM FOR MODELING LOSS OF
RESIDUAL TCA . . . .. 173

APPENDIX C. AREA COUNT DATA SET FOR STATISTICAL ANALYSIS
OF WATER EXTRACTS OF GASOLINE . .. 179

REFERENCES . . . . . 187

BIOGRAPHICAL SKETCH . . . . .. 194















Abstract of Dissertation Presented to the Graduate School
of the University of Florida in Partial Fulfillment of the
Requirements for the Degree of Doctor of Philosophy


BEHAVIOR OF PARTIALLY MISCIBLE ORGANIC COMPOUNDS
IN SIMULATED GROUND WATER SYSTEMS



By

Patricia V. Cline

August 1988

Chairman: Joseph J. Delfino
Major Department: Environmental Engineering Sciences

Serious ground water contamination problems result from

leaks or spills of organic liquids which are partially

miscible in water. Two important categories of these

liquids include low molecular weight chlorinated solvents

and gasoline.

l,l,l-Trichloroethane (TCA) abiotically degrades in

water forming approximately 17-25% l,l-dichloroethene (1,1-

DCE) via an elimination reaction. The substitution product

is acetic acid. The Arrhenius activation energy is 119 +/-

3 kj/mol with an Arrhenius factor of 2 X 1013 s-1, which

results in an estimated half-life for the degradation at

250C of 10.2 months.

Brominated analogs of TCA hydrolyze 11-13 times faster








than TCA. As the number of bromines increase, the percent

of elimination products increases.

These geminal trihalides degrade by a unimolecular

mechanism (E1/SN1). The rate coefficient for TCA

degradation in buffered water at elevated temperature is

approximately six times greater than hydrolysis of 1-

chloropropane (SN2 mechanism) and more than 100 times

greater than l,l-dichloroethane. In the presence of sodium

thiosulfate, the l-chloropropane degradation rate increased

by more than a factor of 100, l,l-dichloroethane rate by 22

and TCA degradation by approximately two.

Halogenated ethenes are stable at various temperatures

and reaction conditions. Trichloroethene degrades in

alkaline solution at elevated temperature.

l,l,l-Trichloroethane and 1,1-DCE form a near ideal

solution in the solvent phase. The solubility of 1,1-DCE at

24C is 3200 mg/l and the solubility of TCA is approximately

1580 mg/l.

The range of concentrations for major components of

gasoline which partition into water was determined for 65

gasoline samples. Benzene concentrations in the water

extracts ranged from 12.3-130 mg/l and toluene

concentrations ranged from 23-185 mg/l.

Fuel/water partition coefficients of seven major

aromatic constituents were measured for 31 gasoline types

and showed a standard deviation of 10-30%. These









coefficients were highly correlated with the pure component

solubilities.

Chemometric techniques were applied to 20 peaks

measured in the aqueous extracts of the 65 gasolines.

Bivariate plots and principal component analyses show

selected brands have distinguishing equilibrium

concentrations, but complete separation of brands was not

observed.


vii















INTRODUCTION


Liquids organic compounds ar frequent causes of ground

water contamination. Nonaqueous-phase liquids (NAPL) fall

into two broad categories based on their migration patterns

upon reaching ground water. Mineral oils, including crude

oils as well as various refined products like gasoline, are

less dense than water and move vertically through the

unsaturated (vadose) zone and tend to spread laterally upon

reaching the water table. The majority of spills involving

organic fluids which contaminate ground water result from

this group of compounds (Schwille, 1984).

In many industrialized countries, serious threats to

ground water supplies result from low molecular weight

chlorinated solvents. These anthropogenic substances are

more dense than water and vertical rather than lateral

movement dominates upon reaching the water table. The more

common solvents detected in ground water include 1,1,1-

trichloroethane (TCA), trichloroethene (TCE),

tetrachloroethene or perchloroethene (PCE), and various

dichloroethene isomers. In addition to common usage, the

high frequency of detection is attributed to the compounds'

high mobility and relatively high resistance to degradation.








2

Decreases in the concentration of contaminants measured

in environmental samples can occur as a result of various

attenuation mechanisms. These include biodegradation,

volatilization, photooxidation, and dispersion. In the

subsurface, losses from pathways like photoxidation are not

important. Other pathways like volatilization occur at

rates slower than those measured from exposed surfaces.

Aerobic biodegradation can occur in the subsurface providing

adequate oxygen and nutrients are available and that the

contaminants are not present in concentrations which are

toxic for microorganisms.

The major objectives of this research include

determining fuel/water partitioning patterns and measuring

chemical degradation rates to aid in the interpretation of

data from contaminated ground water sites. Field

investigations of sites contaminated by gasoline or

chlorinated solvents typically analyze and report the

presence of constituents which are regulated by the state or

federal government (e.g. priority pollutants). These

components are emphasized in my research.

Many chlorinated organic compounds will degrade in

water by hydrolysis or elimination mechanisms. Due to the

extended residence times of organic pollutants in ground

water, this typically slow abiotic degradation within months

or years can be a significant attenuation mechanism. The

focus of my research on the halogenated solvents is on the











transformation processes, and the factors which affect

reaction pathway and the rates of degradation.

Gasoline is a complex mixture of hydrocarbons. Ground

water contamination by gasoline is characterized by elevated

concentrations of the more water-soluble constituents. The

focus of my research on gasoline hydrocarbons is on the

distribution or partitioning of various components of the

gasoline mixtures into ground water and the variability in

the equilibrium concentrations of major constituents.



Chemistry of Alkyl Halides

Abiotic transformation has been reported for TCA, TCE

and PCE, with less work reported on the dichloroethene

isomers. My research reevaluates previous studies and

further examines the chemistry of these compounds.

Mechanisms are evaluated to aid in predicting behavior of

alkyl halides in complex subsurface environments which can

catalyze reactions, lead to the formation of complexes, or

provide localized microenvironments of variable pH or redox

potential.

Evidence of the importance of the abiotic

transformation of l,l,l-trichloroethane (TCA) has been

presented (Cline et al., 1986). l,l-Dichloroethene or

vinylidene chloride (1,1-DCE) was one of the five most

frequently detected volatile organic compounds in finished

drinking water supplies, other than trihalomethanes,










according to a survey by the US Environmental Protection

Agency (Westrick et al., 1984). Vinylidene chloride (1,1-

DCE) is a highly reactive, flammable liquid which is

primarily used in the production of copolymers with vinyl

chloride or acrylonitrile. Emissions occur during

manufacturing, shipping and production; however, these

emissions represent less than 1% of the total 1,1-DCE

produced (Environmental Protection Agency, 1985). The

common occurrence of this compound as a ground water

contaminant cannot be entirely explained by its production

and usage patterns.

One source of 1,1-DCE develops during the abiotic

degradation of 1,1,1-trichloroethane (TCA). The production

of TCA is more than three times the production of 1,1-DCE,

and unlike 1,1-DCE, it is an end-use product indicating that

emission to the environment is essentially equivalent to the

production (Environmental Protection Agency, 1985). The

presence of 1,1-DCE is typically associated with the

presence of other alkyl halides. Since 1,1-DCE is more

toxic than TCA, the conversion to 1,1-DCE in ground water

can increase the toxicity of the water supply.

The association of 1,1-DCE with TCA can be seen more

dramatically in field data from sites which show high levels

of chlorinated solvents in ground water. A summary of

volatile organic compounds (VOC's) in Arizona's ground water

(Graf, 1986) states that, of the six most commonly detected











VOC's, only three (1,1,1-trichloroethane (TCA),

trichloroethene (TCE), and tetrachloroethene (PCE)) are used

in large quantities at the industrial facilities. The

presence of 1,2-dichloroethene isomers and 1,1-

dichloroethane, particularly with frequent detections of

vinyl chloride, suggest anaerobic biodegradation (Parsons

and Lage, 1985; Bouwer and McCarty, 1983). Selected

locations show very high levels of 1,1-DCE in association

with TCA, and frequently little evidence of biodegradation

(Table 1). The primary source of 1,1-DCE at these locations

appears to be the chemical degradation of TCA, prompting

questions as to the rate of formation of the 1,1-DCE and its

stability in ground water.



Table 1. Maximum Concentrations (jg/L) of VOC's Detected
at Selected Sites in Arizona (Graf, 1986)

Site TCA 1,1-DCE TCE 1,2-DCE

1 630 3320 13000 20
2 490 1320 9
3 9800 10400 410 933
4 98 206 139 106


Two products are formed during the abiotic degradation

of TCA. The elimination product is 1,1-DCE, while the

substitution or hydrolysis product is acetic acid (Figure

1). Previous research (Cline, 1987) described the rate of

degradation of TCA and formation of 1,1-DCE in dilute buffer

solutions (pH 4-10) at temperatures from 27 to 700C.








|CI
C= C
I CI


pl
H3C-CQ'+ + CI
Cl
C1


1,1 -Dichloroethene


ELIMINATION PATHWAY


1,1,,1 -Trichloroethane


OH
H3C-C-CI
CI


H3C-C-CI H3C-C-OH

Acetic Acid


SUBSTITUTION PATHWAY


Figure 1. Abiotic degradation pathways for 1,1,1-trichloroethane.


Cl
CI

H3C-C-CI
CI
Cl








7

Transformation processes are most evident in field data

when the degradation products accumulate and are analyzed

and reported, as shown for TCA. The slow degradation of

priority pollutants to products which are not analyzed or

reported (alcohols, aldehydes, or acids which are not

regulated substances) is not as easily characterized in

field investigations. This may occur during degradation of

chlorinated ethenes. The common occurance of TCE and PCE,

as well as the formation of dichloroethene isomers during

degradation, suggest additional study of pathways of the

chlorinated ethenes.

The relative importance of anaerobic biodegradation

versus chemical degradation on a site (Table 1) may be

inferred by observations of the amount of the biodegradation

product of TCE (cis-1,2-DCE) or of TCA (l,l-dichloroethane)

as compared to the chemical degradation product of TCA, 1,1-

DCE. Specific site conditions can affect the relative rates

of these attenuation mechanisms. Abiotic degradation rates

increase as the ground water temperature increases.

Biodegradation rates may be influenced by many factors

including presence of other organic, redox potential,

oxygen concentration, and nutrients.

The specific objectives of my research are to examine

the degradation rate and pathways for halogenated ethanes

and ethenes and determine factors which may affect these

processes.










Gasoline Partitioning

Gasoline contamination of ground water has become a

major environmental concern. Documented cases of

contamination from underground storage tanks (Florida

Department of Environmental Regulation, 1985) have prompted

enactment of additional legislation, the "State Underground

Petroleum Environmental Response Act of 1986" (SUPER Act),

to protect the ground water and surface waters of the state

of Florida. The SUPER act was designed to maximize ground

water protection, encourage early detection, reporting, and

clean-up of leaking underground storage tanks.

Issues relating to the behavior of gasoline components

in ground water are diverse and complex. Gasoline itself is

a complex mixture of hydrocarbons and some of the factors

which affect the concentration of these constituents in the

subsurface environment (vadose zone and ground water)

include solubility, biodegradability, volatility, soil

sorptive capacity, and dilution.

Components of gasoline may undergo abiotic chemical or

photochemical oxidations through free radical formation.

Thermal degradation is negligible at environmental

temperatures below 800C. Since free radicals are limited in

the subsurface environment, chemical degradation is not

expected to play a significant role there (Bossert and

Bartha, 1984).








9

Aerobic biodegradation will be an important attenuation

mechanism provided that sufficient oxygen and nutrients are

present, and these components typically become limiting

after a spill or leak. Attempts to stimulate aerobic

biodegradation of underground petroleum need to remedy both

nutrient and oxygen deficiencies. In addition, hydrocarbons

in the C5-C9 range (which are typical of gasoline) have

relatively high solvent-type membrane toxicity which will

reduce the number of microorganisms and therefore, decrease

the amount of biodegradation following a gasoline spill

(Bossert and Bartha, 1984).

Sites which have been contaminated by gasoline spills

occasionally report results of the analysis of the "floating

layer." Recovery wells to remove the residual organic

liquid are typically installed as an early remediation

measure. Ground water is typically analyzed for benzene,

toluene, and the xylenes (BTX) and more recently for the

oxygenated gasoline additive methyl tertiary butyl ether

(MTBE).

Concentrations of the BTX or oxygenated constituents

will vary spatially and temporally. At the source, changes

in relative concentrations of hydrocarbon components occur

through weathering, primarily volatilization and

solubilization of the liquid residual organic constituents,

resulting in increasing concentrations of the least mobile

constituents. Compounds detected in ground water











downgradient from the spill occur as a result of transport

from the source, and therefore show higher concentrations of

the more mobile constituents.

The downgradient aqueous concentrations are dependent

on the initial partitioning of the gasoline components into

water at the source. The presence of the residual

hydrocarbon will dominate the partitioning process, with

soils playing an increasing role as the residual hydrocarbon

is depleted. Field data are complex to interpret. This is

due to many factors, including site heterogeneities, well

construction and sampling variables, and lack of detailed

information which can provide estimates of the rates of

partitioning and transport. However, patterns resulting

from physical processes, i.e. partitioning and transport,

may be observed. In Table 2 are summarized the highest

concentrations of BTX components measured in monitoring

wells at various gasoline spill/leak sites in Florida.


Table 2. Maximum concentrations (mg/L) of BTX components
in monitoring wells at selected gasoline
contamination sites in Florida.

County Benzene Toluene Xylenes

Hillsborough 24 64 16
11 46 15
Volusia 10 28 11
8 46 9
Desoto 0.8 60 9



These concentrations are similar to those measured in

laboratory gasoline-water partitioning experiments in this








11

study in spite of differences which exist in the age of the

spills and various physical and biological factors. The

time component for the weathering of gasoline at the source

is dependent on many site-specific factors. Even the

relative contributions of volatilization and solubilization

will depend on conditions like the depth of the water table

at the time of the spill and subsequent water table

fluctuations.

A simplification of the complex problem of determining

patterns of gasoline constituent concentrations following a

spill is to initially focus on the partitioning of gasoline

components from the fuel to water. This allows estimations

of equilibrium concentrations of different components from a

fresh.spill in contact with water. Different brands and

grades of gasolines may then be evaluated to determine if

differences among the source types are distinguishable, and

how differences in composition affect the partitioning

behavior.

The major objectives of the gasoline study include

determination of the variability in the fuel/water partition

coefficients for aromatic constituents. Factors which may

affect the partitioning (concentration, cosolvents) will be

evaluated. Chemometric analyses on hydrocarbon components

present in the aqueous solution in equilibrium with gasoline

will be performed to evaluate similarities and differences

in various brands and grades of gasolines.















MATERIALS AND METHODS

Alkvl Halides

Reagent grade chemicals (Fisher Scientific) were used

to prepare buffers and standard solutions. Phosphate

solutions (0.05 M) were prepared at pH 4.5, 7.0 and 8.5 by

mixing stock solutions and monitoring the pH with a Fisher

Accumet model 230A pH meter. Solutions of 0.05 M potassium

dihydrogen phosphate and 0.05 M potassium hydrogen phosphate

were prepared using distilled deionized water. Equal molar

volumes were used for the pH 7.0 buffer. The phosphate

solutions at pH 4.5 (potassium dihydrogen phosphate) and pH

8.5 (potassium hydrogen phosphate) required minor pH

adjustment using 0.05 M phosphoric acid or potassium

hydroxide solutions.

Stock standard solutions of TCA and 1,1-DCE were

prepared in methanol at concentrations of approximately 1

mg/mL. Working standards were prepared by spiking

approximately 5 pL of the stock standard solution into 10 mL

of distilled deionized water. Aliquots of 100-500 pL of.the

working standards were used to prepare standard curves for

the response of the gas chromatograph (GC) to the

concentration of analyte.










Seawater samples were obtained from the coastal

Atlantic Ocean near Ormond Beach, Florida. Samples were

filtered and subsequently handled similar to the phosphate

solutions.

Ground water samples from two monitoring wells were

obtained from a site in Orlando, Florida, which had been

contaminated by chlorinated solvents. These samples were

purged to remove existing solvents and interfering

substances, then filter (10 pm) sterilized.

Approximately 6.6 mL of the phosphate solutions,

seawater or distilled deionized water were added to 5 mL

(nominal volume) glass ampules (Wheaton Scientific). The

ampules were plugged with cotton and autoclaved for 15

minutes at 1210C.

These ampules were then aseptically spiked with 10 pL

of the stock solution of TCA in methanol and flame sealed

using a Model 524PS sealing unit manufactured by O.I.

Corporation. Final concentrations were approximately 1-3

mg/l. Approximately 0.5 to 1 mL of air space was present in

the ampules after sealing.

Ampules were incubated at 280C (Precision Scientific

Model 6) and at 370C (Precision Scientific Model 4).

Experiments at higher temperatures were performed in a

Magna-Whirl constant temperature water bath (Blue M).

Samples were analyzed using a purge and trap device

(Tekmar LSC-2), interfaced with a Perkin Elmer Model 8410 GC










with flame ionization detector (FID) which employed a 30 m

J&W DB-I, 0.53 mm i.d. wide bore capillary column with a 3

pm stationary film thickness. The temperature program

included a 10 minute hold time at 300C and temperature

ramping of 5C/min to 800C. The helium flow was 2.5 mL/min.

Selected analyses were performed by gas chromatography/mass

spectrometry (GC/MS) for quantification and confirmation of

the formation of 1,1-DCE.

The brominated analogs of TCA were not commercially

available. These compounds were synthesized according to

the protocol described by Stengle and Taylor (1970). Two

hundred and fifty milliliters of carbon disulfide (CS2) were

added to a 500 mL, 3-neck flask that was saturated with HBr

vapors at OOC. Excess vapors were trapped over aqueous KOH.

Five milliliters of TCA were added. Five grams of aluminum

bromide (AlBr3) were added to 100 mL anhydrous CS2, placed

in a dropping funnel, and gradually added to the TCA/CS2/HBr

solution over a period of one hour.

This solution was extracted with ice water made basic

with ammonium hydroxide. The solvent was then removed by

distillation and the residue was filtered. An aliquot of

the mixture was added to methanol and spiked into ampules

containing water. Analysis by purge and trap GC showed two

primary peaks and a secondary peak. The major peaks were

determined by GC/MS to be tribromoethane and

dibromochloroethane. A smaller peak was shown to be










bromodichloroethane. Trichloroethane was below detection

levels ( <30 pg/L )in these analyses.

Some of the spiked ampules were heated for a few hours

to determine if halogenated ethenes would be formed, and if

so, to subsequently determine their corresponding retention

times. Two major peaks were identified by GC/MS to be 1,1-

dibromoethene and l-bromo-l-chloroethene. A sample GC

chromatogram containing reactants and products is shown in

Figure 2, with mass spectra of TBA and DBE in Figure 3.

The same analytical conditions were used for the

brominated compounds as were used for TCA, although the

final temperature was slightly increased.

Pure standards of these compounds were not available

for quantification. The degradation rate was determined

directly from the decrease in area counts, since the

response of the external standard remained consistent during

the time of the experiments. However, determination of

molar concentrations was required to determine the percent

transformation to the elimination product.

The response on an FID is generally related to the

number of carbons and can be affected by functional groups.

To determine if the molar response on the FID was affected

by the type of halogen on the molecule (bromine or

chlorine), I examined the response of the trihalomethane

series for which standards were available (Table 3). The

molar response on the FID was the same for this series of


















FID Response


7.88
10. 1 .
II .15

14.?3


4.87



12.30


17...I


.f-9.47
21.3!

5. 04
J- ,28.11


Figure 2. Sample chromatogram of partially degraded geminal
trihalides.


Compound

1,1,1-Tribromoethane
l,l-Dibromo-l-chloroethane
l-Bromo-l,l-dichloroethane

1,1-Dibromoethene
1-Bromo-l-chloroethene
1,1-Dichloroethene


Retention Time

25.93
23.24
19.47

17.33
12.30
7.80


a t-: r PGN

W .


3. 93









17
73S RCT. TIHC: 18.55 TOT AIJV4D- 132141. BASE PK/MrUtli: 136.9 34530.
U
C 100 26.1
a TEA
S TBA
c 80


los
S40

S105
4 0
"j 20
( 93 10 172 251 266
( D) O 1 , i- r ,,
32 80 100 120 140 60 18 200 22 240 260
p m/z *

1MLPVYC/TBA/KCL*e0i1nt4T ,45-450,2:0~OP.24HOV8 7,unD ~ 13S66,~ I 31
DiCS,30M,IUn,8030eS-2SO,S.2D,3SR,SCRYO 874 SCaNS ( 874 SCANS, 15.88 MIMS)
S1.0 MAftS RACCE: 44.0, 269.8 TOTAL 3SUNDs 323S689.

DEE

TBA


SS 110 16S 220 276 331 384 439 494 550 65S 660 715 77 82S


2* 29 RET. TIME: 10.2 TOT GEUII[- 199?422. EFSE PK/ACiUUI': 10S.O/ 46640.

S100 23.4
Sov DBE
*0


4-


S1-,,
I *" o I- -

m/z











Figure 3. Total ion chromatogram (center) for partially
degraded geminal trihalide mixture, with mass spectra for
1,1,1-tribromoethane and 1,1-dibromoethene.








18

compounds. Therefore, the molar response factor for TCA was

used to quantify the ethanes containing bromine and the

molar response factor for 1,1-DCE used to quantify the

brominated ethenes.


Table 3. Relative response


Trihalomethane

1 Chloroform
2 Bromoform
3 Chloroform
4 Bromoform
5 Bromodichloromethane
6 Dibromochloromethane
7 Bromodichloromethane
8 Dibromochloromethane


e of trihalomethanes on

Area
ng nmoles Counts

616 5.15 24.18
924 3.65 15.19
924 7.73 43.98
1386 5.48 22.02
502 3.06 18.47
386 1.85 9.89
1255 7.65 38.79
965 4.63 21.56

Average
Std. Dev.
Rel. Dev.


Response Factor, nmoles/area counts.




Gasoline

Analyses for gasoline constituents were also performed

by GC/FID, using a Perkin-Elmer Model 8410 gas chromatograph

with a 30 m wide bore capillary column (J&W, DB-1) having a

3 pm film thickness. The neat gasoline samples were

analyzed by direct injection of 0.05 pL of the fuel.

Gasoline components dissolved in water were determined by

sparging volatiles from water using a Tekmar LSC-2 Purge and

Trap instrument interfaced to the Perkin-Elmer GC. The

temperature program for both neat gasolines and water


GC/FID


rf*

0.21
0.24
0.18
0.25
0.17
0.19
0.20
0.21

0.21
0.03
13.5%








19

extracts included a 13-minute hold time at 350C, temperature

ramping of 30C/min to 90C, then 5C/min to 2000C. The

helium flow rate was 3.0 mL/min.

Between August and December 1986, subsamples of

gasolines were obtained from the Department of Agriculture

and Consumer Services (DACS) Petroleum Laboratory in

Tallahassee, Florida. These samples were originally

collected by field inspectors and shipped for analysis to

assess compliance with ASTM guidelines and represent various

terminals in northern and central Florida. These samples

represented both summer and winter blends. Subsamples were

collected into 40 mL VOA screw cap vials with Teflon lined

septa and stored on ice prior to analysis.

Local samples were also collected from selected gas

stations in Gainesville. Samples were obtained from the

pump in gasoline safety containers, then a subsample was

transferred to a VOA vial and cooled.

Procedures for evaluating the partitioning of gasoline

into the aqueous phase were reported by Coleman (1984) and

Brookman et al. (1985a). Brookman et al. (1985a) measured

concentrations of aromatic compounds in the aqueous phase

with varying rotation contact times and found a maximum

concentration after two hours. Samples were then

centrifuged to separate the two phases. Coleman et al.

(1984) determined that a rotation contact time of 30 minutes

and an equilibration period of approximately 1 hour produced








20

consistent results and that longer periods had little effect

on the final concentrations.

Saturated, equilibrated solutions of neat gasolines in

contact with distilled, deionized, organic-free water were

prepared. Two mL of gasoline were added to 40 mL water in

VOA vials having Teflon septa. Samples were mixed on a

rotating disk apparatus for 30 minutes at room temperature

(generally 21-230C). The vials then sat undisturbed for one

hour, in an inverted position. Each separated water phase

was removed through the septum at the bottom of the VOA

bottle using a 5 mL syringe. A separate needle was inserted

to allow air to enter the vial so that a vacuum did not form

preventing withdrawal of the water.

Triplicate samples of each water phase were then sealed

in 2 mL crimp-seal vials and refrigerated until the GC

analysis was performed, typically within 2 days. Replicate

extractions, and replicate analyses of extracts were

performed for quality control.

Some overlap or incomplete peak resolution occurred in

the early eluting compounds for both the neat gasoline

samples and the water extracts. Enhancement of the more

water soluble components occurred following aqueous

extraction, making it easier to identify compounds like

benzene and MTBE in the water extract. Toluene was easily

identified in both the neat and water fractions.










When the objective of comparing gasoline samples

involved identification and quantitation of MTBE, analysis

of the water extract provided the most straightforward

interpretation. Although MTBE may be present in gasoline in

quantities approaching 11%, it was more commonly present at

about 5%. MTBE has a lower FID response than the

hydrocarbons, and eluted early in the chromatogram where

several other components also eluted. In samples that did

not contain MTBE, hydrocarbon peaks were present at lower

concentrations at MTBE's retention time. Since MTBE has a

much greater water solubility than these other constituents,

the relative proportion of MTBE to hydrocarbons was

increased in the water extract.















DEGRADATION OF ALKYL HALIDES


Introduction

In this section the degradation kinetics for 1,1,1-

trichloroethane (TCA) and other 1,1,1-trihaloethanes will be

presented and discussed. These compounds degrade in water

forming both elimination and substitution products.

Specific experiments were performed to determine the

mechanism of this reaction and to describe factors which may

effect the rate or pathway of the degradation.

Mechanisms of hydrolysis/elimination have been studied

for many years and numerous reviews, textbook chapters and

empirical concepts have been developed to describe the

chemical degradation of alkyl halides in water. The

following review provides the framework for subsequent

discussions of alkyl halide structure and reaction

mechanisms where specific examples will be presented. The

information was synthesized from several sources (March,

1985; Carey and Sundberg, 1984; Mabey and Mill, 1978;

Bentley and Schleyer, 1977).

Classical SN1, SN2, El and E2 mechanisms have been

defined as early as 1933 (Figure 4). The distinction

between SN1 and SN2 is whether or not the nucleophilic









Unimolecular Mechanisms


I I



-C-C+ + OH-
I I


I I
-C-C+ + X
I I


I I
-C-C-OH
I I


ESh:I


Step 1.



Step 2.



Step 2.


I I-'
-C-C+

H


Bimolecular Mechanisms



OH- + -C-X D- HO C X-- HO-C- + XSN2
i I I


-C-C-
OHI
SH
OH-


C=C
/


+ H20 +


x- E2


Figure 4. Classical substitution and elimination reaction
mechanisms for degradation of alkyl halides in water.


C C
/C = C










attack at the alpha carbon (carbon containing the halogen

leaving group) occurs before the transition state in the

rate determining step, not the extent to which the bond to

the leaving group is broken. Clear cut differences in

substitution reaction mechanisms are apparent in many

reactions. In practice, there is a spectrum of SN2

mechanisms involving varying amounts of nucleophilic attack,

with SN1 being the limiting case where nucleophilic attack

does not occur before the transition state of the rate

determining step.

Unimolecular (SN1 or El) processes are favored by

systems that form stable carbocationsI. A classic example

would be the hydrolysis of t-butyl bromide. The more polar

the solvent, the faster the reaction. An increase in ionic

strength will typically increase the reaction rate, unless

the anion is the leaving group ion (common ion effect). The

reaction is independent of the concentration of nucleophile.

The classic SN2 case occurs in molecules with low

steric hindrance and low carbocation stability. Simple

primary halides react by the SN2 mechanism, while secondary

halides react by an SN2 or intermediate mechanism. Solvent



1 For years these were called "carbonium ions".
Recently, it was determined that the term "carbonium ions"
more accurately refers to pentacoordinated positive ions
(e.g. CH5+) and the more typical positive ion intermediates
(R3C ) are "carbenium ions". The term "carbocation"
includes either type and is generally used to describe any
of these intermediates (March, 1985, p. 141-142).










polarity has less effect on the reaction rate than is

observed for SN1 reactions, but the rate is more sensitive

to changes in concentration or strength of nucleophiles.

The E2 reaction occurs when base attacks the hydrogen

at the carbon adjacent to the carbon containing the leaving

group (beta carbon). This reaction occurs at higher pH and

is more rapid for molecules containing a more acidic

hydrogen.


Degradation of 1.1.1-Trichloroethane (TCA)

The abiotic degradation of TCA was the subject of my

master's thesis (Cline, 1987) which included a detailed

discussion of related degradation studies and illustrations

of the first order decay of TCA in aqueous solution.

Additional data were collected subsequent to those studies.

This included additional concentration measurements in long

term degradation studies and measurements of rate

coefficients in additional matrices. In this section, a

concise comprehensive summary of these data are presented.

A brief synopsis of previous degradation studies of TCA

which have been reported in the literature is summarized

here. Dilling et al. (1975) performed reactivity studies on

selected chlorinated solvents, including TCA. Estimated

rate coefficients were based on four measurements over a

period of one year for each of two sets of reaction ampules;

one set was maintained in the laboratory and a second set

kept outdoors in Midland, Michigan. The same estimated rate











was reported for each experiment, with half-lives of

approximately six months. Reaction products were not

measured.

The hydrolysis of TCA in seawater was reported by

Pearson and McConnell (1975). A half-life of 39 weeks (9

months) was estimated for TCA at 100C with the predominant

reaction being dehydrochlorination to 1,1-DCE. Walraevens

et al. (1974) examined the degradation of TCA in 0.5, 1.0

and 2.0 M sodium hydroxide solutions. The elimination

reaction was not observed, and sodium acetate was shown by

infrared analysis to be the sole reaction product. The

elimination product, 1,1-DCE, was assumed to be stable under

all experimental conditions.

Vogel and McCarty (1987) monitored the degradation of

TCA and formation of 1,1-DCE in water at pH 7 and a

temperature of 20C. The TCA half-life at 200C was

estimated to be between 2.8 and 19 years. Haag and Mill

(1988) report approximately 22% conversion of TCA to the

elimination product, with an extrapolated half-life of 350

days (11.5 months) at 250C.

Degradation experiments were performed at various

temperatures and in different sample matrices. The results

of these experiments are summarized in Table 4. First order

degradation kinetics were observed (Figure 5) in the

data as verified by plotting In [TCA] versus time. Linear

regression analyses were performed on each data set. All

















Table 4. Summary of TCA Degradation Rates
and Product Formation


Temp.
oC


Matrix


108 k
s-1


ke/k
%


70 pH 4
pH 5
pH 7
pH 10
GW1
GW2


62 pH 13

53 pH 4.5
pH 7.0
pH 7.0
pH 8.5
Seawater
DW

39 pH 4.5
pH 7.0
pH 8.5

28 pH 4.5
pH 7.0
pH 8.5


1390 +/- 85
1530 +/- 90
1410 +/-100
1400 +/- 95
1480 +/- 90
1400 +/- 80


565 +/- 35

140 +/- 12
140 +/- 15
144 +/- 20
145 +/- 16
155 +/- 18
133 +/- 14

25 +/- 1.2
24 +/- 1.1
24 +/- 1.2

4.4 +/- 0.2
3.9 +/- 0.2
4.2 +/- 0.2


26 +/- 1

25 +/- 1
26 +/- 2




38 +/- 1

25 +/- 2
24 +/- 2
24 +/- 2
25 +/- 2
"25
23 +/- 3

19 +/- 1
22 +/- 1
17 +/- 1

23 +/- 2
19 +/- 2
21 +/- 2


DW, Distilled organic free water
GW, Ground water matrix


#obs







8.0
7.8
7.6-
.7.4 i
7.2 -
S7.0
6.8 "" --
6.6-
6.4
6.2
6.0 ,
0 100 200 300
Time (days)

350

300

250-

S200

L 150-

100-

50
50-
0 i i -O- -- -- --
0 0.2 0.4 0.6 0.8 1
1 e-kt


Figure 5. First order kinetic data for the degradation of
1,1,1-trichloroethane at 280C and pH 4.5, with the
corresponding data for the formation of the elimination
product, 1,1-dichloroethene.









29

rate constants were based on reactions showing a minimum of

75% degradation of the initial concentration of TCA.

Statistical analyses were performed to assess if the

slopes measured at any given temperature were significantly

different, thus determining the extent to which the sample

matrix, or pH affected the rate constant. The reaction

rates in the buffer solutions (pH 4.5, 7 and 8.5) were not

significantly affected by pH (p < 0.01). In addition, the

rates measured in ground water matrices at 70C (GW1, GW2)

were not significantly different from rates measured in the

buffer solutions at the same temperature.

The spiking solutions typically were prepared with

methanol, which resulted in approximately 0.1% methanol in

the final solution. Separate experiments were conducted

without the use of methanol with no apparent affect on the

rates. The use of methanol decreased the variability in

concentrations observed among ampules, apparently due to the

decreased volatility of TCA in the methanol spiking

solution.

Reaction rates at 530C in seawater, distilled deionized

water and 0.05 M phosphate buffer solutions showed that the

ionic matrix affected the rate of reaction. The fastest

rate was observed for seawater, while the rate in distilled

deionized water (DW) was 14% lower and those in the buffer

solutions were approximately 10% lower. The rates measured

in the distilled deionized water and the buffer solutions











were not significantly different; however, the rate in the

seawater matrix was higher than these at the p<0.01 level.

The 10-14% increase in reaction rate observed in the

seawater matrix at this temperature may be due to the

catalytic influence of some component of that matrix, or to

the increase of ionized species concentration in the

solution.

The relationship between the rate coefficient, k, and

temperature is expressed by the Arrhenius equation,

In k In A EA/RT, where EA is the Arrhenius activation

energy, R is the gas constant, T is the temperature and A is

the Arrhenius pre-exponential factor. The plot of the data

from this and other studies is shown in Figure 6. The plot

includes rates for a variety of matrices including seawater

and sodium hydroxide solutions. Since two products were

formed, the degradation process was complex, but the overall

linearity of the Arrhenius plot implies that a single rate-

determining step is involved in the degradation. Based on

these results, an activation energy of 119+/-3 kJ/mol and an

Arrhenius (A) factor of 2.0x1013 s'1 were calculated.

Extrapolated rate constants and estimated half-lives are

shown in Table 5.

Table 5. Extrapolated Half-Lives for the Degradation of TCA

Temperature (C) Half-life (years)

15 4.5 +/- 0.8
20 2.0 +/- 0.3
25 0.85 +/- 0.13












-10 This dissertation.
v Pearson and McConnell, 1975.
x Vogel and McCarty, 1987.

-12 o Haag and Mill, 1988.
O Dilling et al., 1975.
+ Wolroevens et al., 1974.
-14
Io
O0

S-16-
C-
-J v
-18-



-20



-22 ,
2.8 3 3.2 3.4 3.6

1000/T (OK)

Figure 6. Arrhenius plot for the abiotic degradation of 1,1,1-trichloroethane.









32

Included in the Arrhenius plot are the degradation rate

coefficient for TCA in a pH 13 buffer and also the rate

coefficients calculated by Walraevens et al. (1974) for the

sodium hydroxide solutions. The rates for these high pH

solutions were within the confidence interval for the

regression line, indicating the reaction rate was not

significantly accelerated in alkaline media. The lack of

change in the rate in the presence of a high concentration

of a strong nucleophile (i.e. OH') suggested that the

reaction with the nucleophile occurs after the rate

determining step, characteristic of SN1 reactions.

Similarly, the increase in base strength did not shift the

elimination to an E2 mechanism through a large rate increase

and/or increase in formation of the elimination product.

The rate data which exceeded the confidence interval of

the regression line (Figure 6) were from studies (Vogel and

McCarty, 1987; Pearson and McConnell, 1975) which estimated

the rates of the slow reactions with less than 50%

degradation of the parent compound occurring. Rate

constants calculated for low conversion are more variable

than rates established based on higher amounts of conversion

(Levenspiel, 1972, p. 85). The strong linear Arrhenius

relationship between temperature and rate observed between

25 and 800C, regardless of sample matrix, suggests that

reaction rates at temperatures below 250C can be estimated

by extrapolation.










The elimination product, 1,1-DCE, was measured to

establish the factors which influenced the reaction pathway

(substitution versus elimination). Degradation of 1,1-DCE

was observed only at very high pH and even under those

conditions the rate was slow compared with the degradation

of TCA. Therefore, the ratio of the rate for elimination

(ke) to the total rate of degradation (k) was estimated by

plotting the concentration of 1,1-DCE versus (l-e-kt) where

t is time. The slope of the line equals ([TCA]o (ke/k)),

where (TCA]o is the concentration of TCA at time zero.

This calculation required an estimate for the starting

concentration of TCA. For most experiments, multiple

analyses were performed for the estimate of the initial

concentration. Other authors (eg. Vogel and McCarty, 1987)

have used the intercept in the regression analysis of the

degradation, and this value was used as the estimate of

initial concentrations in this study.

Increases in pH and/or temperature theoretically favor

elimination over substitution. The elimination pathway

(Table 4) ranged from 17 to 38% of the total degradation

rate of TCA. Higher temperatures showed slightly more

transformation to 1,1-DCE over the temperature range

evaluated in these experiments. The percent of TCA

degradation due to elimination was not affected by matrix in

the pH range of 4.5 to 8.5. Seawater had no apparent effect

on the relative proportion of products. The highest percent












elimination pathway was measured in the strongest sodium

hydroxide (pH 13) solution.

Qualitative observations (GC and GC/MS) of TCA

degradation at approximately 600C in 0.5, 1.0 and 2.0 molar

sodium hydroxide solutions, showed the presence of 1,1-DCE,

and separate experiments indicated that 1,1-DCE also slowly

degraded under those conditions. These findings contradict

the results reported by Walraevens et al. (1974) in which

1,1-DCE was not detected in TCA degradation experiments at

high pH. This may be due to differences in analytical

methods, or the slow degradation of 1,1-DCE under their

reaction conditions.


Degradation of Brominated Ethanes

The degradation rates of brominated versus chlorinated

1,1,1-trihaloethanes were compared to provide insight into

the mechanisms and overall behavior of these compounds.

Since bromine is a better "leaving group" than chlorine,

brominated compounds typically degrade faster than their

chlorinated counterparts. In reviewing hydrolysis

degradation processes, Mabey and Mill (1978) concluded that

Br is more reactive than Cl by a factor of 5 to 10.

In a search of Chemical Abstracts, fewer than 20

references were reported for the brominated analog of TCA,

1,1,1-tribromoethane (TBA). Most of the papers addressed

spectra and bond energy studies, while no information on the

hydrolysis of this compound was reported.











Brominated analogs of TCA were not commercially

available. Therefore, TBA was synthesized according to the

methods reported by Stengle and Taylor (1970). The

procedure for the synthesis of 1,1,1-tribromoethane (TBA)

produced a mixture of brominated analogs of TCA. The

primary components were TBA and l,l-dibromo-l-chloroethane

(DBCA), while smaller quantities of l-bromo-l,l-

dichloroethane (BDCA) were present. Kinetic data for

abiotic degradation of TBA and DBCA were measured for

several temperatures while data for BDCA were obtained in

only selected experiments conducted at higher overall

concentrations. Compound structures are illustrated in

Figure 7. The elimination pathway involved loss of HBr to

form the corresponding alkene, the dominant elimination

product was the ethene formed by loss of a bromine. The

substitution pathway forms acetic acid.

Initial degradation experiments involving the

synthesized brominated mixture were conducted in reagent

grade (Milli-Q) water to obtain preliminary data on the

transformation process. Subsequent experiments were

conducted in buffer solutions at pH 4, 7, and 10. The

results of these experiments are summarized in Table 6.

First-order kinetics of degradation were observed, as

were also seen for TCA. Rate constants were calculated

from the linear regression analysis of the plots of the













H Br
HC-CBr
H Br


1,1,1-Tribromoethane (TBA)


H Br
HC-CBr
H CI


1,1-Dibromo-1 -chloroethane (DBCA)


H Br
HC-CCI
H CI


1 -Bromo-1.1 -dichloroethone (BDCA)


/Br
C=C
/ \Br


1,1-Dibromoethene (DBE)


S /Br
C=C


1-Bromo-1 -chloroethene (BCE)


C=C
/ \CI


1,1-Dichloroethene (DCE)


Figure 7. Brominated analogs of 1,1,1-trichloroethane and
corresponding elimination products. Since bromine is a
better leaving group than chlorine, the predominant pathway
is elimination of HBr.














Table 6. Summary of Brominated Compound Degradation
Rate Coefficients and Product Formation


Tem


p.


Matrix Compound


20 DW
20 pH 4
20 pH 7
20 pH 10
28 DW
30 pH 4
30 pH 7
30 pH 10
37 1 M KC1
65 DW
65 Na2S203

20 DW
20 pH 4
20 pH 7
20 pH 10
28 DW
30 pH 4
30 pH 7
30 pH 10
37 1 M KC1
65 DW
65 Na2S203


65 DW


TBA
TBA
TBA
TBA
TBA
TBA
TBA
TBA
TBA
TBA
TBA

DBCA
DBCA
DBCA
DBCA
DBCA
DBCA
DBCA
DBCA
DBCA
DBCA
DBCA


BDCA


108 k
s-1

14
10
9
11
42
65
64
71
492
7700
13000

17
14
11
15
51
81
69
73
492
7860
14000


5350


ke/k


50.9
60.7
58.5
61.8
64.1
61.8
56.3
63.1
38.0
51.6
60.1

35.6
33.9
32.5
33.4
45.6
39.8
31.9
38.5
24.4
33.8
40.6


29.6


DW, Distilled organic free water
Na2S203, 1 M Sodium thiosulfate


Extrapolated Half-Lives for Degradation of TBA and DBCA


TBA


Temp.


15
20
25


108 k
s-1

5.54
12.83
28.92


T 1/2
Days

145
62
28


108 k
s-1

6.90
15.56
34.14


DBCA


T 1/2
Days

116
52
23











natural log of the concentrations versus time. All rate

constants were based on reactions showing a minimum of 75%

degradation.

The results of the degradation of TBA, DBCA and BDCA at

65C are illustrated in Figure 8. The differences in

slopes for the degradation of these compounds were not

statistically significant indicating that the rate

determining step was similar for each compound.

The formation of products (Figure 9) was calculated as

discussed previously for the formation of 1,1-DCE. The

percent elimination (ke/k) was the slope of the regression

line divided by the initial concentration of the parent

product. The smaller slope for 1,1-DCE, and its lower

maximum concentration, was a function of both lower initial

concentration of reactant (BDCA) and lower percent of BDCA

degradation which occurred through the elimination pathway.

The Arrhenius plot for TCA as determined in this study

is compared in Figure 10 with that of the brominated

compounds, TBA and DBCA. The Arrhenius plot for the two

brominated compounds was represented by a single regression

line. The regression line for TCA was essentially parallel

to that of the brominated compounds. The Arrhenius

activation energy (EA) for all of these compounds was almost

identical, since EA is a function of the slope of this line.

The rate of degradation of TCA at 25C was

approximately a factor of 11 to 13 times slower than for the


























0 2 4 6
Time (Hours)
Figure 8. First order kinetic data for the abiotic
degradation of TBA, DBCA and BDCA in water at 650C.


45-
40-
35-
30-
25-
20-
15-
10-


5

0


0.2


0.4


0.6


0.8


-kt
1-e-

Figure 9. Formation of the elimination products (BCE, DBE,
DCE) in water at 650C from the abiotic degradation of the
corresponding 1,1,1-trihaloethanes.


+ BCE :
* DBE ,
* DCE ++


+
+ +
+



S0 0











+ TBA
DBCA
0 TCA
-10




Y -15




-20




2.8 3 3.2 3.4 3.6

1000/T (OK)

Figure 10. Arrhenius plot for the abiotic degradation of 1,1,1-trihaloethanes.










brominated analogs. The observed rate constants for the

various pH values were not significantly different, as was

also observed for the degradation of TCA.

Experiments were performed to determine the reaction

mechanism for these 1,1,1-trihaloethanes. First, the

degradation experiment was conducted in a 1 molar sodium

thiosulfate solution. Sodium thiosulfate is a much stronger

nucleophile than water or hydroxide (Swain and Scott, 1953)

and a dramatic increase in degradation rate in this solution

is indicative of an SN2 reaction in which the nucleophile is

directly involved in the rate determining step. The

degradation rates measured for the brominated 1,1,1-

trihaloethanes increased less than a factor of 2 in the

thiosulfate solution, which may be attributed to the

increased ionic strength of the solution. The percent

elimination was also unaffected by this sample matrix.

To further characterize the mechanism for these

degradation reactions, the brominated geminal trihalides

were placed in a 1 M KC1 solution at 370C. High ionic

strength solutions generally increase the rate of SN1 or El

reactions. When a common ion is present, the rate of the

reverse reaction is enhanced. In the presence of high

concentrations of chloride, chloride may be exchanged for

bromide when an ion pair forms. If BDCA forms an ion pair

and chloride is exchanged, TCA will be formed (Figure 11)

providing evidence of a carbocation intermediate. Even










though TCA will degrade, it is more stable than the

brominated compounds and it may accumulate to detectable

levels.

The BDCA compound had the lowest concentration in the

mixture of the three geminal trihalides in reaction

solution, and TCA concentration was less than 40 ug/l.

After three days of incubation at 370C, the concentration of

TCA rose to approximately 200 ug/l. This was a minor

pathway (less than 5% of the BDCA degraded forming

detectable TCA) in the overall degradation process. 1,1,1-

Trichloroethane was not detected in other sample matrices

during the degradation experiments of the brominated

compounds, indicating that its presence in this solution was

a result of the reverse reaction of carbocation with the

chloride in the solution.

Increasing the extent of bromination increased the

percent of the degradation resulting in the elimination

product (Table 6). The proportion of the total degradation

which resulted in elimination for BDCA at 650C was within

the error estimate for the percent elimination of TCA at

elevated temperatures, and both of these parent compounds

produced 1,1-DCE. The highest percent elimination was

observed for TBA which formed approximately 60% 1,1-

dibromoethene (Figure 12). This may be due to an increase

in steric hindrance in carbocations containing bromine

rather than chlorine, slowing the substitution pathway.









H Cl
HC-CCI
H CI


Exchange
Reaction


H CI
HC-CBr
H Cl

BDCA


H CI
HC-C+
H CI


Ion Pair


Br-


/ El
----v


SN1


HC-C
H \OH
OH


Acetic Acid


Figure 11. Reaction pathways for BDCA in 1 M KC1 solution. The exchange reaction of Cl
with the ion pair forms TCA, which degrades more slowly than the brominated compound.


H
\
HC
H








100

90-

80-

70-

60-

S50-

2 40-

30-

20-

10-

0


Figure 12. Comparison of the percent of the elimination pathway for 1,1,1-trihaloethanes.


TBA


DBCA


BDCA


I~1


TCA










These experiments provided evidence that the abiotic

degradation of 1,1,1-trihaloethanes occurred by SN1/El

rather than SN2 and E2 mechanisms. The trihaloethanes

containing one or more bromine atoms degraded at similar

rates, approximately a factor of 11-13 faster than TCA,

reflecting that bromine was a better leaving group. As the

number of bromines present on the trihaloethanes increased,

the percent of the degradation occurring through the

elimination pathway increased.


Degradation of Halogenated Ethenes

One of the primary objectives of examining the behavior

of halogenated ethenes was to provide an accurate evaluation

of their formation and stability during degradation of the

corresponding ethanes. The literature provided some

evidence that slow degradation of these ethenes may occur at

a rate of interest for ground water studies.

Supporting the possibility of degradation, Dilling et

al. (1975) reported half-lives for the abiotic degradation

of trichloroethene (TCE) of 10.7 months (0.002 day-1) and

for tetrachloroethene (PCE) of 9.9 months at 250C.

Molecular oxygen was present and the degradation rates were

suggested to result from oxidation as well as hydrolysis.

In this often referenced work, it was suggested that

mechanisms of degradation at lower temperatures may differ

from rates extrapolated from studies at higher temperatures.










In a study of hydrolytic decomposition by Pearson and

McConnell (1975), volatilization was extrapolated to zero

and a degradation half-life for TCE of 30 months was

estimated.

Roberts (1985) examined field evidence for the

degradation of various chlorinated organic and estimated

rate constants for both TCE and PCE of approximately 0.003

day-1, which may be due to a variety of factors including

sorption and dilution.

Wilson et al. (1985) studied the aerobic degradation of

TCE, PCE and other compounds in actual aquifer materials

from two sites in Oklahoma and Louisiana. No detectable

biodegradation of these compounds was observed under the

experimental conditions. Since degradation was noted in

autoclaved samples, the authors postulated that TCE and PCE

degradation was likely due to abiotic processes with rates

similar to those reported by Dilling et al. (1975).

The dehydrochlorination reaction of TCE occurs under

basic conditions and generates dichloroacetylene and

hydrogen chloride. This reaction of TCE with base is

spontaneous at room temperature and was responsible for

dichloroacetylene intoxication observed in patients inhaling

TCE-containing air in closed systems equipped with alkali

absorbers (Environmental Protection Agency, 1979).

Dichloroacetylene was detected in the gas phase above

aqueous alkaline solutions with pH 11 to 13 and upon











incubation with moderately alkaline material such as

concrete (Greim et al., 1984). They concluded

dehydrohalogenation can occur under these relatively mild

conditions resulting in toxicity from exposure to the

dichloroacetylene.

Many substitution and addition reactions of TCE have

been carried out in the presence of base. What initially

appeared to be a direct substitution reaction may in fact

have been multistep processes involving intermediates like

carbanions, chloroacetylenes, or carbenes. Rappaport (1969)

reviewed the mechanisms for nucleophilic vinylic

substitution processes in alkaline solutions at elevated

temperatures.

Mechanisms may differ for chemical studies performed

under extreme conditions of temperature and high pH compared

to reactions occurring under more typical environmental

conditions. The possibility of slow nucleophilic attack in

aqueous solution was considered because March (1985) reports

that although vinyl halides are generally considered

resistant to nucleophilic attack, the presence of electron-

withdrawing groups like halogen lower the electron density

of the double bond enhancing nucleophilic substitution or

addition reactions.

In ground water, even very slow degradation may be an

important attenuation mechanism. Since environmental

studies report slow degradation of TCE or PCE in water and











the chemical studies show presence of electron withdrawing

groups like chlorine increases the susceptibility of an

olefin to nucleophilic attack, experiments to evaluate

possible reactions were performed.

References to possible hydrolysis reactions of 1,1-DCE

or its brominated analogs were not found upon review of the

literature. The 1,1-dihaloethenes would be less susceptible

to nucleophilic attack than TCE since fewer electron

withdrawing groups are present. The pure compounds however,

are very reactive and polymerize readily. Their

reactivities in dilute aqueous solution have not been

examined.

The focus of my research with halogenated ethenes was

to examine the stability of these compounds in relatively

dilute aqueous solutions and to determine their

susceptibility to nucleophilic attack. Autooxidation or

other reactions of the pure liquid compounds which may be

present in the vadose zone following a spill could occur,

but these reactions are not addressed here.

There were two major purposes for the examination of

the degradation behavior of halogenated ethenes. First, the

stability of the ethene products formed during the

transformation of the geminal trihalides needed to be

determined to accurately describe the kinetics of the

appearance of these elimination products. Secondly,

previous studies which indicated that halogenated ethenes












like trichloroethene (TCE) and tetrachloroethene (PCE) may

undergo slow abiotic degradation in water at room

temperature with a half-life of less than one year were

reevaluated. The question of possible nucleophilic attack

by water, hydroxide ion, or other nucleophiles must be

addressed to understand the stability of these commonly

detected ground water contaminants.

The stability of 1,1-DCE was evaluated in experiments

that were performed concurrently with the evaluation of TCA

degradation. In the buffer solutions, seawater, and

distilled deionized water, no significant degradation of

1,1-DCE occurred during the course of the evaluation of the

degradation of TCA.

The formation of ethenes containing bromine was

monitored during the degradation studies of the brominated

ethanes, and their concentrations were continually monitored

for some time after the ethane degradation was completed.

Trichloroethene was studied in separate experiments

performed at various temperatures selected to repeat the

experiments conducted by Dilling et al. (1975). In addition

to buffer solutions, one set of ampules was prepared with a

nutrient solution which was not autoclaved, and to which

ground water known to show biological activity was added.

This was done to determine if any degradation which might

have occurred during the long term studies could have been

due to biological activity.










A summary of the results of these experiments is

presented in Table 7. No significant degradation of these

compounds was found in the experimental matrices during the

indicated reaction times, as evidenced by the slopes of In C

vs time which were not significantly different from zero.

The overall coefficient of variation for the observations is

similar to values obtained for simple replicate analyses.

Experiments were also conducted to evaluate the overall

behavior of these compounds under more rigorous conditions.

The literature indicated that halogenated ethenes such as

TCE can undergo elimination to form chloroacetylenes at

elevated pH (Rappaport, 1969). This reaction was verified

by using GC/MS to confirm the formation of dichloroacetylene

from TCE and also chloroacetylene from 1,1-DCE by analysis

of the headspace vapor above an alkaline (1 M NaOH) aqueous

solution of the halogenated ethene which was warmed to

approximately 600C.

The rate of degradation of components in a mixture of

1,1-DCE, TCE and PCE in sodium hydroxide solutions was

examined at 600C (Table 8). These were the matrices used by

Walraevens et al. (1974) in their examination of the

degradation of TCA, wherein they did not observe formation

of 1,1-DCE. One objective was to establish if the

elimination product was stable under their reaction

conditions.











Summary of Experimental Conditions for which
Halogenated Ethenes were Stable.


Time No.
Cmpd. Days Obs.


Temp.
oC


27
27
27
27
37
65

27
27
27
27
37
65


BCE
BCE
BCE
BCE
BCE
BCE

DBE
DBE
DBE
DBE
DBE
DBE

DCE
DCE
DCE
DCE
DCE
DCE
DCE
DCE
DCE
DCE
DCE
DCE
DCE
DCE

TCE
TCE
TCE
TCE
TCE
TCE
TCE
TCE
TCE
TCE


Average
Concentration
mg/L


4.2
4.2
4.2
88
80
392


32
32
28
120
110
690


160
160
160
54
126
6

160
160
160
54
126
6

274
386
386
386
386
386
386
15
15
15
14
14
14
6

274
386
386
386
386
386
386
14
14
14


C.V. Matrix


7%
9%
6%
4%
4%
11%

4%
7%
10%
2%
11%
12%

7%
13%
14%
12%
9%
8%
10%
16%
6%
9%
12%
9%
14%
12%

4%
9%
10%
8%
12%
9%
16%
8%
9%
9%


pH 10
pH 7
pH 4
DW
DW
Thio

pH 10
pH 4
pH 7
DW
DW
Thio


Nutrient
pH 4
pH 7
pH 8.5
pH 4
pH 7
pH 8.5
DW
pH 7
Seawater
pH 4
pH 7
pH 8.5
Thio

Nutrient
pH 4
pH 7
pH 8.5
pH 4
pH 7
pH 8.5
pH 4
pH 7
pH 8.5


DW Distilled organic free water
Thio 1 M Sodium thiosulfate solution


Table 7.


2.7
1.1
1.1
1.1
2.4
2.4
2.4
2.1
2.2
2.3
2.3
2.3
2.3
.4

3.0
1.2
1.2
1.2
1.7
1.7
1.7
1.7
1.7
1.7













Table 8. Second Order Degradation Rates (1 mole-1 hr-1)
of Halogenated Ethenes at 600C
in Sodium Hydroxide Solutions

NaOH TCE 1,1-DCE PCE
Concentration
0.1 M 0.6 0.02 nd
0.5 M 0.28 0.01 nd
1.0 M 0.17 0.01 nd
2.0 M 0.12 0.004 nd

nd no significant degradation occurred after 260 hours.



The rate of degradation of TCE was the greatest among

the tested compounds due to the presence of an acidic

hydrogen (a hydrogen present on a carbon containing a

halogen). The elimination reaction was also an available

pathway for the degradation of 1,1-DCE, although the rate of

degradation was approximately 30 times slower than for TCE

in all solutions except for the 1.0 M NaOH.

Tetrachloroethene (PCE) did not degrade since the

dehydrohalogenation reaction could not occur, and apparently

conditions were not favorable for an addition process.

For environmental applications, there are concerns with

the mildest conditions (temperature and pH) which may still

result in degradation of these compounds. The pathway for

the degradation of ethenes at elevated temperature could

differ from reactions at lower temperatures where the

elimination reaction would be less favorable and a possible

addition reaction could occur instead. Therefore, TCE was








53

incubated at 200C in a solution at pH 12.5. No degradation

was observed during four months of incubation (Table 7).

The experiments demonstrate the resistance of the

halogenated ethenes to degradation in dilute aqueous

solution. Reports of the degradation of these compounds

with half-lives of less than 1 year appear to represent a

process other than abiotic degradation in water. In the

same way as Dilling et al. (1975), my experiments were

conducted in sealed ampules containing a headspace, however

degradation was not observed as reported in their study. I

believe their results may be a result of analytical error.

The half-lives for each experiment were based on four

measurements. The results showed a chemically diverse group

of compounds had similar decreases in concentration and

temperature had little effect on these decreases. A

possible explanation for these results would be a decrease

in instrument response over the year of the study.

The halogenated ethenes generally showed very little

degradation, with the exception of the rapid degradation of

TCE at high pH and temperature. It appears that any

degradation of these compounds in aqueous solution which

occurs, does so under rather extreme conditions and is not

expected to be a dominant process.


Structure/Rate Relationships of Alkyl Halides

In the previous sections degradation patterns and

kinetics were evaluated for various 1,1,1-trihaloethanes in












aqueous solution. A broader perspective on hydrolysis /

elimination reactions can be obtained by comparisons with

other haloalkanes reported in the literature. The

objectives are

1. To compare degradation rates measured for

trihaloethanes of other simple alkyl halides which react by

an SN1/E1 mechanism.

2. To compare degradation rates of trihaloethanes with

other geminal trihalides reported in the literature to

determine structure/activity relationships with changes in

the substituents on the beta carbon, and describe shifts in

mechanisms which may occur for these trihalides.

3. To compare degradation rates and pathways of 1-

chloropropane and l,l-dichloroethane with TCA to show the

effects of increasing number of chlorines on the alpha

carbon.

The classic reaction mechanisms for substitution and

elimination reactions are SN1, SN2, El and E2, as previously

discussed. The presence of various functional groups can

effect the rate and pathway of degradation of an alkyl

halide. For example, rates of hydrolysis are greater for

alkyl halides containing Br rather than for Cl by a factor

of 5 to 10. The rates also increase as the alkyl group goes

from primary to secondary to tertiary in the ratio of

1:10:1000 for chloride. Allyl groups enhance the rate of

hydrolysis of a primary halide by a factor of 5 to 100,










while benzyl groups enhance the rate by a factor of 50.

(Mabey and Mill, 1978)

The formation of stabilized carbocations by electron

donation from the non-bonded electron pairs of halogens

adjacent to the cationic carbon center have been reported

(Olah, 1974). The stabilizing effect was enhanced when two

or even three electron-donating heteroatoms coordinate with

the electron-deficient carbon atom as illustrated in Figure

13. Specific examples, designated as "chlorocarbenium

ions" by Olah (1974), have been identified and are

illustrated in Figure 14.

Simple SN1/E1 Reactions

My data suggested 1,1,1-trihaloethanes form carbocation

intermediates. The intermediate would contain two halogens

and one methyl group. The observed rates and pathways are

compared (Table 9) to compounds containing two methyl groups

and one halogen (2,2-dihalopropanes) and three methyl groups

(t-butyl chloride).

Degradation of tertiary halides like t-butyl chloride

occurs with a carbocation intermediate and these compounds

are resistant to bimolecular nucleophilic displacement. The

half-life for the aqueous degradation of t-butyl chloride is

approximately 23 seconds at 25C with about 19% of the

degradation occurring through the elimination pathway. The

carbocation intermediate is stabilized by the three methyl












+
RC
2


+
x-C c


- x


+
X =


R C
2


c -x


Figure 13. Stabilization of carbocations by halogen (Olah, 1974).


x


zX


X c
I







H

C -CC12
H3
;-< ^ cl


HC


-I
I


- CH3


CH3

-CCI2


-~cI2


CH3


Figure 14. Examples of "chlorocarbenium ions" (Olah, 1974)


3


HC-
3C


~Ilr?


















Table 9. Summary of Degradation Rate Coefficients and
Pathways for Tertiary and Secondary Halides


Compound


t-Butyl Bromide
t-Butyl Chloride

2-Bromo-2-chloropropane
2,2-Dibromopropane
2,2-Dichloropropane

1,1,1-Tribromoethane
l,l-Dibromo-l-chloroethane
1,1,1-Trichloroethane

2-Bromopropane
2-Chloropropane


k (sec-1)
2500


2.98x10-2

1.78x10-4
4.62x10-5
9.09x10-6

2.89x10-7
3.41x10-7
2.62x10-8

3.82x10-6
2.11x10-7


ke/kt
%


Reference


100
100
100


References:
1. March, 1985.
2. Queen and Robertson, 1966.
3. This dissertation.










groups. The rate coefficient at 25C is approximately 106

faster than for TCA.

Queen and Robertson (1966) examined the hydrolysis of

2,2-dihalopropanes. These compounds form carbocation

intermediates with two methyl and one halogen group. The

rate coefficients for the degradation of 2,2-dihalopropanes

are intermediate between t-butyl chloride and the 1,1,1-

trihaloethanes. The mechanism was reported to be SN1/El

based on results of experiments with deuterated gem-

dihalides. The degradation rates of these compounds were

10-50 times higher than of the corresponding secondary

halides (e.g., 2-chloropropane).

The degradation rates were affected by the leaving

group, bromine or chlorine. Also, the structure and

stability of the resulting carbocation affected the rate and

pathway (elimination and/or substitution) of the reaction.

Since bromine was a better leaving group than chlorine,

there was a rate increase when bromine was present as

compared to the corresponding chlorinated compound. 2,2-

Dibromopropane degraded 19 times faster than 2,2-

dichloropropane, while 2-bromo-2-chloropropane degraded 5

times faster than the dichloro compound (Queen and

Robertson, 1966). The 1,1,1-trihaloethanes containing

bromine degraded 11-13 times faster than TCA.

Rates were also increased as the number of methyl

groups present on the carbocation increased. The t-butyl










chloride degraded approximately 3000 times faster than 2,2-

dichloropropane and 106 faster than TCA.

There were two major differences between my results and

those reported by Queen and Robertson (1966). First, they

reported a rate nearly four times higher for 2-bromo-2-

chloropropane than for 2,2-dibromopropane, while the rate

coefficients I measured for the trihaloethanes containing at

least one bromine were approximately equal (within 20%).

Secondly, they report only formation of the elimination

product for all 2,2-dihaloethanes, while the percent

elimination in my experiments was a function of the number

of bromines and was always less than 60%. The percent

elimination for t-butyl chloride was less than the value

obtained for the trihaloethanes.

The effect of alpha halogen is complex, "combining a

negative inductive effect and an electron-releasing

resonance effect" (Queen and Robertson, 1966, p. 1364).

Based on my results and the results for t-butyl chloride,

elimination was not expected as the primary pathway nor the

large difference in rates observed for the two

dihalopropanes which contained a bromine. The rate data

were determined for the dihalopropanes by a conductance

method. Extraction of the products of solvolysis of 2,2-

dibromoethane with CC14 and analysis by vapor phase

chromatography (GC) and nmr showed 2-bromopropene was the

only product in other than trace amounts. It may be that










the substitution product, acetone, would not have

partitioned and been measured using that analytical

protocol.

Comparisons of Geminal Trihalides

A number of compounds in the literature contain a

geminal trihalide group (R-CX3), and many of these compounds

have environmental implications. My experiments on 1,1,1-

trihaloethanes indicated that the -CX3 group was sterically

hindered and resistant to attack by an SN2 mechanism, and

that the halogens could help to stabilize the formation of a

carbocation. The overall rate of degradation of other

geminal trihalides will increase if R also stabilizes the

carbocation. If the beta carbon contains an acidic hydrogen

the mechanism may shift to E2 at elevated pH.

A summary of degradation rates (expressed as reaction

half-lives) of various geminal trihalides is presented in

Table 10. The simplest compounds, trihalomethanes, were

very resistant to hydrolysis. The R- consists only of

hydrogen, which was inadequate to stabilize a carbocation.

The mechanism for this degradation has been determined to be

a base catalyzed process with a carbanion intermediate

(Hine, 1950). The extremely low reactivity also suggests

that steric hindrance may prevent SN2 attack.

By contrast alpha,alpha,alpha-trichlorotoluene has a

half-life of 19 seconds at 250C, which corresponds to a rate

of a factor of 106 greater than for TCA. Therefore, the












Table 10. Half-lives for Abiotic Degradation of
Geminal Trihalides


COMPOUND


STRUCTURE


HALF-LIFE
250C


REFERENCE


Chloroform


Bromoform


CHC13


CHBr3


1,1,1-Trichloroethane CH3CC13


1,1,1-Tribromoethane CH3CBr3


DDT


cI C H CC[3

2


Methoxychlor


3500 yr


690 yr



10.2 mo


1 mo


Mabey and Mill,
1978

Mabey and Mill,
1978


This dissertation


This dissertation


12 yr Wolfe et al., 1977
(pH5, 270C)


1 yr Wolfe et al., 1977
(pH5, 270C)


(CH 0 CH CCI3

2


a,a,a-Trichlorotoluene / CC a
Cd/ 3


19 s Lyman et al., 1982










rate increase was much greater than the factor of 50

reported by Mabey and Mill (1978).

Quemeneur et al. (1971) determined that tri-chloro

compounds of the type p-RC6H4-CC13 (R is OMe, Me, H, Cl, or

NO2) were hydrolyzed in neutral or acidic medium via a

cationic transition state for all types of R substituents.

The hydrolysis of the p-substituted alpha,alpha-

dichlorotoluenes reacted via a cationic mechanism when R is

an electron-donor, and a bimolecular mechanism when R is an

electron-attracting group. These results also supported the

observation that halogens contributed to the stability of

the carbocation. Monochlorotoluene reacts nearly 3000 times

more slowly by an SN2 mechanism than the trichlorotoluene

reacts by the SN1.

Methoxychlor and DDT are two environmentally important

pesticides which contain a geminal trihalide functional

group. Wolfe et al. (1977) provided an in depth examination

of the degradation of these compounds. There is a beta

hydrogen on each of these compounds. At elevated pH the

degradation rate increased as a function of pH and the

elimination products were dominant, suggesting these

structures were more susceptible to degradation by the E2

mechanism than is TCA. While the elimination product, DDE,

was the major product of DDT hydrolysis even at lower pH,

the major product of methoxychlor at pH 7 was the hydrolysis

product, with minor amounts of the elimination product,










DMDE. The hydrolysis products formed were anisoin and

anisil, which were explained by phenyl group rearrangement

after the formation of the carbocation.

Mochida et al. (1967) showed 1,1,1,2-tetrachloroethane

and pentachloroethanes reacted more slowly than TCA under

lower pH conditions, which indicated that chlorines on the

beta carbon decrease the stability of the carbocation. The

presence of these chlorines on the beta carbon however,

increased the acidity of the hydrogens, with enhanced

degradation rates for the tetra and pentachloroethanes by an

E2 mechanism at elevated pH.

There is considerable evidence that geminal trihalides

can form carbocations in the presence of an appropriate

neighboring group. Subsequent reaction pathways may vary

according to the structure of the carbocation resulting in

elimination, substitution, or rearrangements. An E2

reaction may also occur for compounds containing an acidic

hydrogen on the beta carbon.

Effect of Additional Halogens on the Alpha Carbon

The hydrolysis of a simple primary halide, 1-

chloropropane, was compared with the reactivity of 1,1-

dichloroethane and TCA in experiments I performed at

elevated temperature. As the number of hydrogens on the

alpha carbon decrease, steric hindrance can increase and

result in a shift in reaction mechanism. The experiments

were designed to demonstrate the relative rates of










hydrolysis in aqueous solution, and the response to an

increase in concentration of a strong nucleophile whose

effect would be a function of the mechanism.

Based on the literature, simple primary alkyl halides

like l-chloropropane are expected to degrade by an SN2

mechanism. Therefore, l-chloropropane should show an

increase in degradation rate in the presence of a strong

nucleophile, since the nucleophile is involved in the rate

determining step.

Predicting the degradation rate of l,l-dichloroethane

is more difficult. Secondary chlorides, like isopropyl

chloride, have been shown to degrade more quickly than the

primary alkyl halides, possibly by an intermediate

mechanism. Chloride can contribute somewhat to the

stability of a carbocation, however, it is not as effective

as a methyl group as discussed previously. In addition, the

presence of a halogen can increase the steric hindrance at

the alpha carbon.

Comparisons of the degradation rates of these compounds

were made at elevated temperature (650C) in pH 7 buffer

solution, and in a 1 M thiosulfate solution. In the buffer

solution the degradation of TCA was approximately 6 times

faster than the hydrolysis of 1-chloropropane. Degradation

of 1,1-dichloroethane was less than 6% of the rate of 1-

chloropropane degradation. This rate comparison is

illustrated in Figure 15.



















O

o 2



1 +






0 200 400 600
Time (hours)
Figure 15. Pseudo-first-order kinetic data plots for hydrolytic degradation of TCA,
1-chloropropane, and 1,1-dichloroethane in pH 7 buffer solution at 650C.








67

The degradation of 1-chloropropane was enhanced by more

than a factor of 100 in the thiosulfate solution, 1,1-

dichloroethane degraded approximately 22 times faster, and

TCA degradation rate increased less than a factor of 2. The

differences in rate enhancement among these compounds is

attributed to differences in mechanism. Part of the

increase in rate of degradation of TCA in thiosulfate is

attributed to the increasing ionic strength, and TCA

degradation rate was clearly less affected by the presence

of thiosulfate than the other compounds. The rate

enhancement for l,l-dichloroethane was similar to the type

of rate increase which would be observed for secondary

halides which react by an intermediate mechanism.

The thiosulfate solution was used as a matter of

convenience as a strong nucleophile to assist in

demonstrating how knowledge of mechanism may be necessary in

estimating degradation rates as matrices change. Greatest

changes in rates in the presence of sulfur nucleophiles may

be expected for simple primary alkyl halides, and the least

effect occur with compounds which react via an SN1 or El

mechanism.


Sediment Matrix Effects

There is considerable interest in possible effects of

solid surfaces on rates of hydrolysis. Most hydrolysis

experiments are performed in simple buffered aqueous

solution. Contaminants in the vadose zone or ground water










have considerable contact with a variety of aquifer

materials which could potentially affect degradation rate.

Hydrolysis reactions may be affected by factors like acid or

base catalysis, sorption and ionic strength. Since

compounds which react by different mechanisms may be

impacted differently by these solid surfaces, both 1-

chloropropane and TCA were used in degradation experiments

performed in various matrices.

Catalysis of hydrolysis or elimination reactions of

alkyl halides by saturated aquifer materials has not been

demonstrated. Because high concentrations of 1,1-DCE have

been observed in Florida and Arizona at solvent spill sites

contaminated with TCA, the role of sand or other materials

which may influence the degradation of TCA was evaluated.

The nonbiological degradation of pesticides in the

unsaturated zone was shown to play an important role for a

few groups of pesticides, mainly organophosphates and s-

triazines. Clay mineral surfaces have shown catalytic

activity, correlated to their acid strength. This catalytic

process is most important at low moisture content, and

therefore is more important in the vadose zone than beneath

the water table (Saltzman and Mingelgrin, 1984).

Haag and Mill (1988) did not observe significant

differences in the kinetics or products of TCA in contact

with sediment pore water. Epoxide hydrolysis was










accelerated by a factor of four in sediment as compared to

rates in buffered water.

Mabey and Mill (1978) indicated that acid promotion of

the aqueous hydrolysis of halogenated aliphatic hydrocarbons

has not been observed. March (1985) stated that gem-

dihalides can be hydrolyzed in water with either acid or

basic catalysis to give aldehydes or ketones, although the

strength of acid was not addressed.

In a review of elimination reactions in the presence of

polar catalysts, Noller and Kladnig (1976) stated that

"interaction of X with an acid is probably as indispensable

as the reaction of H with base in liquid-phase elimination

reactions, but this function is probably taken over by the

solvent and is less pronounced than the base promoted

process."

Clarification of interactions with polar surfaces may

provide insight into possible effects of sediments or soil

on reaction rates. Clays, for example, contain polar

surfaces which have been shown to catalyze degradation of

some pesticides (Saltzman and Mingelgrin, 1984).

Noller and Kladnig (1976) illustrated elimination

reaction products were a function of the specific catalyst

with 1,1,2-trichloroethane


Cl H
Cl C1 C2 Cl
H H










as reactant. Basic catalysts (e.g., KOH-SiO2) attack the

most acidic H, that at C1, forming more 1,1- than 1,2-

dichloroethene. Acidic catalysts (e.g., silica-alumina)

attack C1 on Cl because the formation of the carbocation is

facilitated by the other Cl on that carbon resulting in the

formation of much more of the 1,2-dichloroethene isomer.

The choice of catalyst will determine the predominant

product giving selectivity to the reaction.

Mochida et al. (1967) reported that the reactivity of

TCA on solid acids was greater than that for other

chlorinated ethanes (mono-, di-, tri- and tetra- chloro

compounds). On solid bases it was less reactive than

penta-, tetra-, and 1,1,2-tri- chloroethanes. The shift in

reactivity of the ethanes with change in catalyst showed

enhanced ability of TCA to form a carbocation by

accelerating the reaction on an acid surface as compared to

the other chlorinated ethanes. There was also the lack of

an acidic beta hydrogen to permit catalysis by base.

Possible catalysis would be compound- and mechanism-

specific. Degradation experiments were performed on 1-

chloropropane (SN2) and TCA (SN1,E1) at 650C in 5 ml

distilled deionized water, with a final concentration of

approximately 2 mg/l. Separate ampules were prepared with

the addition of 0.4 g bentonite clay, 1 g limestone, 1 g

sand, and 0.2 g silica gel.








71

Similar trends were observed for both compounds (Table

11). The slowest rates relative to water were obtained for

both compounds in the sample containing clay, while the

fastest rates were observed in the sand.

The data generally showed greater variability in the

samples containing the solids as compared to the DW system

(Figures 16 and 17) as evidenced by correlation coefficients

less than 0.99. However, the rates of l-chloropropane

degradation in ampules containing solids differed by less

than 10% of the rate obtained for Milli-Q water.

The relative degradation rates for TCA differed more as

a function of matrix than observed for chloropropane,

however, there was also greater variability as evidenced by

the correlation coefficients. In the case of TCA, the

formation of 1,1-DCE was similar in all matrices suggesting

the ratio of products was not affected by the presence of

these solids.

The relatively small differences in rates measured in

these matrices may be due to a variety of factors including

sorption, however significant surface catalysis was not

observed. For this type of saturated system, the amount of

alkyl halide in contact with the surface would be small.

Differences may be attributed to normal variability and

differences in ionic strength or composition of the aqueous

phase in contact with the solids.













Table 11. Matrix Effects for Degradation Rates of
1-Chloropropane and 1,1,1-Trichloroethane at 700C.
Linear Regression Output for the Plot of in C (ug/1)
vs. Time (hours).

Chloropropane


Regression Output: MQ


Clay Limestone Sand


Silica Gel


Constant (Ci)
Std Err of Y Est
R Squared
No. of Observation
Degrees of Freedom


7.20
0.12
0.99
12
10


7.59
0.26
0.94
9
7


7.44
0.07
0.99
7
5


7.14
0.13
0.99
11
9


X Coeff. (Rate) -0.0102 -0.0094 -0.0098 -0.0113

Std Err of Coef. 0.0004 0.0009 0.0004 0.0004


Relative rate
(MQ 1)


1.00


0.92


0.96


1.11


1.1.1-Trichloroethane


Regression Output:

Constant (Ci)
Std Err of Y Est
R Squared
No. of Observations
Degrees of Freedom


MQ Clay Limestone Silica Gel Sand


5.86
0.11
0.99
11
9


6.26
0.12
0.99
7
5


5.23
0.07
1.00
6
4


5.91
0.18
0.98
7
5


6.01
0.28
0.96
9
7


X Coeff. (Rate)
Std Err of Coef.

Relative rate
(MQ 1)


-0.027 -0.020 -0.038
0.0008 0.0009 0.0010


1.00


0.74


1.38


-0.034 -0.040
0.0020 0.0030


1.25


1.45


7.11
0.23
0.96
12
10

-0.0107

0.0007

1.06

























0 --a
0


0 20


Figure 16. Effect
rate of hydrolytic


40 60 80 100 120 140
Time (hours)


of the presence of
degradation of TCA


solid material
at 650C.


on the


Clay
+ Silica Gel
0 Limestone
x
4 Milli-Q Water
Sand

3- A


2-
0

1 + 0
-_ _


100


200


300


400


Time (hours)

Figure 17. Effect of the presence of solid material on the
rate of hydrolytic degradation of l-chloropropane at 650C.








74

These experiments do suggest that TCA in sand aquifers

may show a slightly increased rate as compared to low ionic

strength buffered water experiments. The rate coefficient,

however, will fall within the error limits for the rate

estimate for the degradation of TCA based on the experiments

in buffered distilled water.














SOLUBILIZATION AND DEGRADATION OF RESIDUAL TCA

A computational model was constructed to describe the

attentuation of TCA beneath the water table in the presence

of multiple phases. This simplified scenario for a TCA

spill considered the chemical transformation of TCA to 1,1-

DCE along with advective transport resulting from ground

water flow, of TCA and 1,1-DCE out of this zone containing

the residual solvent. Biodegradation of TCA in this highly

contaminated zone was considered negligible.

The major objective in developing this model was to

describe the relative concentrations of the major

constituents and how their concentrations may change with

time. These trends are illustrated for various ground water

flow rates, change in initial concentrations, and initial

composition.

Behavior of Residual Solvent

The migration pattern of chlorinated hydrocarbons

following a spill is illustrated in Figure 18. These dense

nonaqueous phase liquids (NAPL) will infiltrate the porous

media, with some of the NAPL retained in residual

concentration. The retention capacity for these NAPL in the

unsaturated zone may range from 5 L m-3 (approximately 12

mL/L of pore space) in highly permeable media to 30-50 L m-3








76







Vadose Zone


.. ..... ................................................................


G r o u n d ... ..... ............................................
W a te r .........................
Flow
----...::.---- II ....i....iKE.....uo4 S

S.of.. e:.Pore.sp e...:.....
::::::::::::::::: ::: :: ::::: :::::::::: :::::::::::::::::::::::T C A
:- : --" ' "," 1 D' '
S.............................................................. ..C




W at.................................................................a t e r
..................:::::..........:::::. Saturated
-..:with Solvent






Chlorinated Solvent Pool
at impermeable layer




Figure 18. Equilibrium model for the attenuation of
residual TCA present beneath the water table.








77

in media of lower permeability (Schwille, 1984). Additional

factors which influence whether the NAPL will reach the

water table include the spilled volume and infiltration

process.

If sufficient volume of dense NAPL reach the water

table, it will sink into the saturated zone and continue to

migrate downward as long as the retention capacity of the

zone is exceeded. Wilson and Conrad (1985) reported

residual hydrocarbon occupying 15-40% of the pore space in

the saturated zone.

Water continues to flow laterally through the water

saturated zone containing residual NAPL. The globules of

NAPL provide a large interface with the water providing a

solution zone, where the initial concentration of a given

component is proportional to its aqueous solubility as

determined by the NAPL composition. These globules are

generally trapped in the larger pore spaces and are being

prevented from entering the smaller pores due to the high

capillary entrance pressure. There is a reduction in

permeability to water where the residual NAPL is present, as

the largest channels become blocked at several places by

discontinuous solvent ganglia. This forces water to flow

around the solvent in fairly thin films and/or be diverted

into the smaller channels whose carrying capacity

(conductivity) is low (Jones, 1985).








78

In laboratory experiments, the initial concentration of

chlorinated solvent was at saturation concentration even

when the layer of sand with residual solvent was thin

(Schwille, 1988). The concentration gradually decreased

until the levels in the water were sufficiently low that

further removal of solvent was slow. At this point

approximately 86% of the residual had been removed.

My model was developed assuming that equilibrium

saturation was maintained, the dissolution of residual

solvent being faster than the degradation or advective

transport of components. Diffusion or hydrodynamic

dispersion was not considered to be a limiting factor in

maintaining equilibrium. The solvent-contaminated zone was

then treated similar to a well-mixed flow reactor.

Interactions of the solutes in the water with the solid

matrix of the saturated zone were considered minimal

providing residual solvent was present; the porous medium

was assumed to provide a matrix in which the residual

solvent was retained.

Once the flow of the NAPL stopped, the subsequent

losses were assumed to occur through degradation or

advection of the compound in the aqueous phase.

Hydrolysis/elimination of TCA occurs much faster in dilute

aqueous solution than would occur for water dissolved in the

TCA solvent phase (Walraevens et al., 1974). Ground water

continues to flow through this zone, although at somewhat










reduced velocities, carrying dissolved components out of

this zone.

Aqueous Phase Concentrations

The quantity of solvent lost each day by advection or

degradation is a function of the concentration of each

component (TCA and 1,1-DCE) in the aqueous phase, which in

turn depends on the composition of the residual NAPL. The

solvent phase may contain TCA and/or 1,1-DCE, or another

solvent which may have been spilled with the TCA.

The distribution of a component between the two liquid

phases can be expressed in terms of fugacity. For ideal

mixtures, the solubility of the solute at any composition is

estimated by multiplying the unit solubility by the mole

fraction of the component in the solvent phase at

equilibrium. Nonideal mixtures form deviations from

linearity. Estimates of aqueous concentrations resulting

from a nonideal solvent mixture requires knowledge of the

activity coefficients at the various mole fraction

compositions. For the simpler ideal case,

[TCA]w x STCA

[DCE]w (l-x) SDCE

x TCAs / (TCAs + DCEs)

where TCAs and DCEs are the number of moles of that compound

in the solvent phase at equilibrium, x is the mole fraction

of TCA in the solvent phase, STCA and SDCE are the pure

component solubilities, and [TCA]w and [DCE]w are the










aqueous phase concentrations at equilibrium. The total

number of moles of TCA in a unit volume of porous media is

the sum of the moles present in the aqueous and solvent

phases.

The model describes changes for TCA spilled on a high

permeability material like sand. As TCA degrades and forms

1,1-DCE, the degradation product partitions into the NAPL

affecting the aqueous phase concentration of TCA (and DCE).

Both the individual solubilities and the solubility of

a mixture of TCA and 1,1-DCE are required in the model and

it was also necessary to assess if mixtures of TCA and 1,1-

DCE deviate significantly from ideality. Literature values

for the solubilities of these constituents vary widely

(Table 12). The solubility data for 1,1-DCE reported by

Lyman (1981), showed as much as a 700% error from a

predicted concentration based on regression relationships.

That estimated concentration is much closer to the

concentrations reported by Verschueren (1977).

Table 12. Solubilities of TCA and 1,1-DCE (mg/L)

Temp (0C TCA 1.1-DCE Source
20 480 400 Pearson and McConnell (1975)
20 4400 2640 Verschueren (1977)
30 1088 3675 Verschueren (1977)
25 273 Lyman (1982)

4 1700 4200 This study.
24 1580 3200 This study.



Measurements (Figure 19) were made on the solubility of

the individual components (TCA and 1,1-DCE) and on the













3000
3000- + + DCE







0
0 +

o i







0 0.2 0.4 0.6 0.8 1.0
Equilibrium Mole Fraction of TCA
(Solvent Phase)
Figure 19. Aqueous solubilities of a binary mixture of TCA and 1,1-DCE as a function of
mole fraction composition in the solvent phase (24 C).










solubility of each with varying compositions of the binary

mixture. Mixtures were at room temperature, approximately

24C.

The pure component solubility of TCA (1580 mg/L or 11.8

mmoles/L) and the solubility of 1,1-DCE in the aqueous phase

(3200 mg/L or 33 mmoles/L) measured at 24C were within the

concentration range listed by Verscheuren (1977) who

reported solubilities at 20 and 300C. This is significantly

higher than solubilities reported by Lyman (1981) and

Pearson and McConnell (1975). The solubility for 1,1-DCE

reported in this dissertation was verified independently by

solubility measurements performed using high performance

liquid chromatography (HPLC) (Linda Lee, University of

Florida, Personal communication, 1988). She measured an

average for the solubility of 1,1-DCE at 24C as 2990 mg/L.

Her report is included in Appendix B.

Verscheuren (1977) reported that the solubility of TCA

at 200C was four times greater than at 3000, a value

approximately three times greater than our result at 240C.

Since the mass lost per unit time from degradation is a

function of aqueous concentration and the first-order

degradation rate coefficient, higher aqueous concentrations

at lower temperatures could compensate for the lower

degradation rate. The solubility of TCA at 40C was measured

to verify this trend. As shown in Table 12, a significant








83

increase in solubility of TCA at lower temperatures was not

observed.

The linearity of the change in solubility with

increasing mole fraction for these two compounds suggested

that 1,1-DCE and TCA form a near-ideal solution in the

solvent phase. Based on these measured data, I assumed that

mole fraction in the solvent phase multiplied by the aqueous

solubility of the pure compound provided a reasonable

estimate of aqueous phase concentration of TCA and 1,1-DCE.

Advection

Loss of TCA from this hypothetical contaminated zone

occurs via advection and degradation, both of which are a

function of the aqueous phase concentration. The relative

importance of these two mechanisms is a function of the flow

velocity advectionn) and the temperature (solubility and

degradation rate). Observations of selected field data

suggest higher concentrations of 1,1-DCE appear in southern

state aquifers where the ground water temperatures are

higher. The model therefore, assigns a temperature of 250C.

The volume of water exchanged through the contaminated

zone is a function of the ground water flow velocity and the

length of the contaminated zone. Fresh water upgradient of

the spill enters the contaminated zone while an equal volume

of water at equilibrium saturation of the contaminants is

displaced. Velocities for the model are expressed as the

per cent of the volume of contaminated water exchanged per







84

day. These values include the "no flow" or "low flow" (0.1%

per day) cases, in which the dominant loss occurs through

degradation. At 0.25% per day, the rate of advection is

comparable to the rate of degradation. Finally, a flow rate

of 0.5% per day represented the case in which the loss of

TCA is primarily due to advection. At flows greater than

0.5% per day the losses would be dominated by the advective

term. These volume exchange rates represent slow flows

and/or very large spill areas. An exchange of 0.5% per day

represents an approximate flow through 5 meters of

contaminated porous media at a rate of 2.5 cm/day.

Degradation Rate

The solubility of TCA affects not only its rate of

advection from the contaminated zone, but also the total

mass of TCA degraded per unit time. The first-order rate

constant at 250C is approximately 0.00226 day-1 as measured

in this study. In a contaminant plume, the half-life for

the degradation of TCA is approximately 10.2 months.

Although the first-order rate coefficient remains constant,

the mass of TCA converted per unit time decreases as the

concentration of TCA in the aqueous phase decreases.

In the model, it was assumed that the TCA concentration

remained at saturation within the zone containing residual

solvent since the TCA that degraded was replaced by

dissolution of the residual solvent. The amount of TCA

degraded per unit time follows zero-order kinetics. The










zero-order rate equals the mass converted per unit time in

the first-order equation as the time increment approaches

zero. This becomes 0.00226 day-1 multiplied by the aqueous

concentration of TCA. A 50% decrease in the solubility

would therefore result in a corresponding 50% decrease in

the mass of TCA degraded per unit time.

Model Parameters and Procedures

Initial conditions for the model include a unit volume

of water (1 liter) in contact with 100 mmoles of TCA. After

equilibrium 11.8 mmoles of TCA will be in the aqueous phase

leaving 88.2 mmoles (approximately 11.8 grams or 8.5 mL) in

the residual solvent phase. The changes in concentration of

TCA or 1,1-DCE in this unit volume are displayed graphically

illustrating the effects of different flow rates, higher

initial mass of TCA, and effect of the presence of an inert

solvent mixed in the residual phase.

Iterative calculations (Appendix C) are made in the

model for advection and degradation in relatively small time

increments, with subsequent reequilibration of the solvent

remaining in the zone of residual contamination. The

residual solvent mass will continue to decrease until at

some point a separate solvent phase does not exist.

Calculations become more difficult (smaller time increments

must be used to attain convergence of the iterative

mathematical solution) and other factors would become more

important as the NAPL is depleted. Therefore, the









86

calculations are stopped when amounts of TCA in the residual

NAPL are less than 10 mmoles. At lower levels of residual

NAPL, the process may become diffusion limited as the NAPL

is trapped in regions of the soil matrix removed from the

aqueous flow. The results of the model are shown in Figures

20-26.

The total mass of TCA in the NAPL showed zero-order

decay with flows from 0.1-0.5% per day (Figure 20). As the

flow rate decreases, slight nonlinearity is observed. This

reflects the slow accumulation of 1,1-DCE in the solvent

phase which begins to decrease the aqueous concentration of

TCA.

The decrease in aqueous concentration of TCA (Figure

21) as the total mass of TCA in the system goes from 100

mmoles to approximately 15 mmoles (slightly in excess of the

solubility) is dependant on the flow. The larger decrease

is observed for the case of no-flow, which results in a 45%

decrease in the aqueous phase concentration after 10 years.

The major reason 1,1-DCE fails to accumulate

significantly in the solvent phase is its higher water

solubility. Having a solubility twice that of TCA, 1,1-DCE

is advected from the zone containing residual solvent more

readily. In the special case of no flow through the system,

1,1-DCE is not advected and begins to accumulate in the

solvent phase affecting the aqueous phase concentration of

TCA. However, since only approximately 20% of the TCA is








100


80


60


40


20


0 2 4 6 8 10


YEARS


Figure 20. Model results: Decrease in total TCA mass in
the residual zone as a function of flow.


C)
Q)
0
F-
E


Q)





H---


YEARS


Figure 21. Model results: Change in aqueous concentration
of TCA as a function of flow.








88

converted to 1,1-DCE, the effect of the accumulation is not

observed until substantial degradation has occurred. If all

the TCA degraded in this closed system, 20 mmoles of 1,1-DCE

would be produced, which is 60% of the pure component

aqueous solubility of 1,1-DCE. Therefore, for the initial

conditions of the model, a residual NAPL will exist only

when excess TCA is present.

A comparison of different initial conditions for a

constant flow (0.25%) is shown in Figure 22. With an

increase in amount of residual TCA, the same zero-order

decay rate is observed, indicating that doubling the amount

of TCA in the solvent phase doubles the time needed for

removal of the residual.

In addition, Figure 22 illustrates the rate of loss of

TCA when the initial 100 mmoles is mixed with another

solvent, a hypothetical mixture in which the mole fraction

of the "inert" compound remains at 0.5 in the solvent phase.

This represents a case where a compound with solubility

similar to TCA (like TCE) is present in the residual. The

presence of this other compound causes a 50% reduction in

the aqueous phase concentration of TCA, and therefore the

rate of loss of TCA, doubling the time to remove the TCA

from the residual phase.

The patterns of change in mass of 1,1-DCE in the

solvent or aqueous phase over time are more complex when

there is advection from the system Figure 23. The aqueous







200
180
160
140
120
100
80
60
40
20
0


YEARS (Flow, 0.25%/Day)


Figure 22. Model results: Change in total mass of TCA as a
function of initial mass of TCA and composition of the
solvent phase.


YEARS (Flow, 0.25%/Day)

Figure 23. Model results: Pattern of 1,1-DCE formation and
advection as 100 mmoles of TCA in the residual zone
degrades.








90

concentration of 1,1-DCE continues to increase for some time

as the mass of 1,1-DCE in the solvent phase begins to

decrease because its mole fraction continues to increase in

the solvent phase.

The total mass of 1,1-DCE in the zone of residual

contamination increased over time, reaching a maximum as the

TCA mass in the solvent phase approached zero. Increasing

the flow rate not only shortened the time in which 1,1-DCE

was accumulating, but decreased the maximum amount of 1,1-

DCE present in that zone. This is true for the aqueous

phase concentrations (Figure 24) and amount in the solvent

phase (Figure 25). The maximum concentration of 1,1-DCE in

the aqueous phase for a flow of 0.5% per day is

approximately 1 mmole/L (100 mg/L) at the point where some

residual phase is still present. The concentration of TCA

at that time is nearly at saturation (approximately 1500

mg/L).

The changes in aqueous concentration of 1,1-DCE for

larger amounts of TCA originally present or in the presence

of an inert solvent as previously discussed, are shown in

Figure 26. The changes in the amount of 1,1-DCE in the

solvent phase is shown in Figure 27. The inert solvent

increases partitioning into the organic phase, keeping the

aqueous concentration low.

The model illustrates factors which affect the time for

removal of a residual phase under varying conditions, and
























2-


0 2 4 6 8
YEARS

Figure 24. Model results: Increase in aqueous
concentration of 1,1-DCE forming from degradation
a function of flow.

7 .


10


of TCA as


YEARS


Figure 25. Model results: Pattern of accumulation of 1,1-
DCE in the solvent phase as TCA degrades.








2.8

2.4

2.0

1.6


1.2

0.8

0.4


0 2 4 6
YEARS (Flow, 0.25%/Day)
Figure 26. Model results: Change in aqueous concentration
of 1,1-DCE as a function of initial mass of TCA and
composition of the solvent phase.
7
---- 100 mmole
-- 200 mmole and
6
100 mmole TCA +
inert organic solvent
5
LU

0 4-


2 3

2-

1

0 ,
0 2 4 6 8 10
YEARS
Figure 27. Model results: Change in total mass of 1,1-DCE
in the residual zone as a function of initial mass of TCA
and composition of the solvent phase.










the different concentrations of 1,1-DCE which would result.

Given a constant initial mass of TCA, the maximum

concentration of 1,1-DCE in the aqueous phase occurs at the

lowest flow rates. For flow rates higher than the 0.5%

volume exchange per day the advective term is dominant and

concentrations of 1,1-DCE in the residual zone remain

negligible.

As long as a residual NAPL is present, aqueous

concentrations are dominated by TCA. Equal concentrations

of TCA and 1,1-DCE in the water from monitoring well data

from various sites would occur according to the model

primarily in the plume of dissolved constituents

downgradient from the residual zone, or in the original

spill area after all residual solvent was dissolved or

degraded. The presence of a low solubility compound in the

solvent phase with the TCA will considerably slow TCA rate

of advection and degradation.

First-order degradation will continue in the ground

water plume downgradient from the source and this process

could be modeled (Kinzelbach, 1985). Evidence of the

formation of 1,1-DCE would support the assignment of a

degradation rate. Assuming similar retardation factors for

TCA and 1,1-DCE, equal concentrations of TCA and 1,1-DCE

would occur after approximately 3 half-lives, approximately

2.5 years at 250C.




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