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1M mu

DEGRADATION OF POLYDIMETHYLSILOXANE AND DIMETHYLSILANEDIOL

IN WATER AND SOIL BY FENTON REAGENT

Hristina Gueneva-Boucheva

A THESIS

Submitted to
Michigan State University
in partial fulfillment of the requirements
for the degree of

MASTER OF SCIENCES

Department of Crop and Soil Science

1999

ABSTRACT

DEGRADATION OF POLYDIMETHYLSILOXANES AND DIMETHYLSILANEDIOL
IN WATER AND SOIL BY FENTON’S REACTION

By

Hristina Gueneva—Boucheva

Silicones fluids consist of polydimethylsiloxanes (PDMS) and are widely used in
industrial and consumer products and enter the environment as components of municipal
sludges applied to the land, or through accidental spills. PDMS derivatives are thought to
be extremely resistant to degradation under ambient environmental conditions. Previous
studies have shown that in soil PDMS degrades to dimethylsilanediol (DMSD) depending
on the environmental conditions. Fenton’s reagent, in which the decomposition of
hydrogen peroxide is catalyzed by iron (II) to form hydroxyl radicals, has recently been
applied to remediation of contaminated soils and wastewaters. Here for the first time the
Fenton reaction was used as a simple and rapid method for PDMS and DMSD
mineralization in water and soil. This reaction successfully breakdown the PDMS and
DMSD at a laboratory conditions. The efficiency of the Fenton reagent to destroy DMSD
was affected by initial H202, iron and DMSD concentrations as well as the oxidation state

of the iron catalysts and iron counterions.

ACKNOWLEDGEMENTS

My research advisor for the past three years has been Professor Stephen Boyd. I
have benefited greatly from his considerable abilities as a scientist and a teacher. I thank
him for his support and guidance during the course of this work.

I also wish to thank Dr. Yucheng Feng who devoted much of her time to getting
me started on my research and for her continuous help and willingness to tech me about
soil sciences and perfect analytical practices.

This work was supported by Dow Corning Inc, Midland, MI

iii

TABLE OF CONTENTS

LIST OF TABLES ................................................................................... vi
LIST OF FIGURES ................................................................................. vii
GENERAL INTRODUCTION .................................................................... 1
Terminology ................................................................................. 2
Environmental occurrence of PDMS ..................................................... 3
Environmental fate of DMSD ............................................................. 9
Fenton chemistry .......................................................................... 11
Application of the Fenton reaction ...................................................... 13
REFERENCES ...................................................................................... 19
CHAPTER I ......................................................................................... 24
INTRODUCTION ........................................................................ 24
MATERIALS AND METHODS ........................................................ 25
Materials ........................................................................... 25
Methods .......................................................................... 26
RESULTS AND DISCUSSION ........................................................ 29
Mineralization of DMSD by Fenton reagent in water ....................... 29
Mineralization of DMSD by Fenton reagent in soil slurries ............... 38
SUMMARY ............................................................................... 47
REFERENCES ............................................................................ 48

CHAPTER H ....................................................................................... 51

INTRODUCTION ............................................................... 5 1
MATERIALS AND METHODS ...................................................... 53
Materials .......................................................................... 53
Methods .......................................................................... 54
RESULTS AND DISCUSSION ....................................................... 55
Degradation of PDMS by Fenton reagent in water ................ 55

Degradation of PDMS by Fenton reagent in soil .................. 64
CONCLUSIONS ......................................................................... 69
REFERENCES ............................................................................ 69

LIST OF TABLES

Table I—l. Characteristics of the Capac A soil ................................................. 26

Table I—2. Kinetic parameters for DMSD degradation in water with differen iron
catalysts ............................................................................................. 34

vi

LIST OF FIGURES

INTRODUCTION
Figure 1. Overview of PDMS environmental mass flow as a result of wastewater

treatment practices. Fraction of sludge disposed by various techniques. Obtained from

EPA survey ( 15) ....................................................................................... 6

CHAPTER I

Figure I-l. Mineralization of DMSD in water by Fenton’s reagent with different iron
catalysts concentrations (lmM DMSD, 440 mM H202) ....................................... 31

Figure 1-2. Mineralization of DMSD in water by Fenton’s reagent with different iron
catalysts (lmM DMSD, 2mM Fe(II)/Fe(III), 440 mM H202) ................................. 32

Figure I-3. Mineralization of DMSD in water by Fenton’s reagent with different H202
concentrations (lmM DMSD, 2 mM iron catalyst) ............................................. 36

Figure 14. Mineralization of DMSD in water by Fenton’s reagent with different DMSD
concentrations (2 mM iron catalyst, 440 mM H202) ............................................ 37

Figure [-5. Mineralization of DMSD in soil by Fenton’s reagent with different iron
catalysts concentrations (10 mM DMSD, 1.5 M H202) ......................................... 39

Figure [-6. Mineralization of DMSD in soil by Fenton’s reagent with different iron
catalysts (10 mg/kg DMSD, 100 mM Fe(II)/Fe(III), 0.625 mM H202) ....................... 41

Figure I-7. Mineralization of DMSD in soil by Fenton’s reagent with different H202
concentrations (10mg/kg DMSD, 100 mM iron catalyst) ...................................... 43

Figure I-8. Mineralization of DMSD in soil by Fenton’s reagent with addition of H202 at
the fifth hour (10 mg/kg DMSD, 100 mM iron (III) catalyst, 0.625 M H202) ............. 44

Figure I-9. Mineralization of DMSD in soil by Fenton’s reagent with different DMSD
concentrations (100 mM iron catalyst, 0.625 M H202) ........................................ 45

Figure I-lO. Mineralization of DMSD in soil by Fenton’s reagent at different solutionzsoil
rates (10 mg/kg DMSD, 100 mM FeCl2, 2 M H202) ........................................... 46

vii

CHAPTER 11

Figure 11-1. Degradation of PDMS in water by Fenton reagent at 66°C and addition of
H202 (0.05 mg/L PDMS, 2 mM FeCl2) .......................................................... 57

Figure II-2. Degradation of PDMS in water by Fenton reagent and FeCl3 and H202 at
room temperature (0.05 mg/L PDMS, 2 mM FeCl2/FeCl3, 440 mM H202) ................ 58

Figure 11-3. Degradation of PDMS in water by Fenton reagent at different temperatures
and addition of H202 at 13 ........................................................................ 59

Figure II—4. Degradation of PDMS in water by Fenton reagent at room temperature and
different PDMS concentrations (0.05 mg/L PDMS, 2 mM FeCl2/FeCl3, 440 mM
H202) ................................................................................................. 60

Figure 11-5. Degradation of PDMS in water by Fenton reagent at different PDMS
concentrations and addition of H202 (2 mM FeCl2, 880 mM H202, and 66°C ) ............ 61

Figure 11-6. Degradation of PDMS in water by Fenton reagent at 66°C and addition of
H202 with and without THF (0.05 mg/L PDMS, 2 mM FeCl2 , 440 mM 202) .............. 62

Figure 11-7. Degradation of PDMS in water by Fenton reagent and by FeCl3 and H202 at
66°C and addition of H202 (0.05 mg/L PDMS, 2 mM FeCl2/FeC13, 440 mM 202) ........ 63

Figure 11—8. Degradation of PDMS in soil by Fenton reagent at room temperature and
different PDMS concentrations (100 mM/kg FeCl2, 0.7 M/kg
H202) .................................................................................. . ............... 66

Figure II-9. Degradation of PDMS in soil by Fenton reagent at different temperatures and
different H202 concentrations (10 mg/kg PDMS, 100 mM FeCl2, 440 mM H202). . . . . ....67

Figure 11-10. Degradation of PDMS in soil by Fenton reagent at 66°C and addition of
different concentrations of H202 (10 mg/kg PDMS, 100 mM FeCl2) ........................ 68

viii

INTRODUCTION

Silicone is used in a broad range of food, medical, personal care and household
applications. It enters the environment mainly through land disposal of wastewater
treatment plants sludges. Although uncommon, accidental spills “or improper usage of
Silicones could contrubute to contamination of soils and waters. With this in mind it is
prudent to explore the environmental fate of silicone products and possible ways to
manage accidental releases of Silicones.

Fenton’s reagent has been applied to the remediation of contaminated soil (1) and

wastewater (2). The term Fenton reagent or Fenton reaction is used to describe an iron

(ID/hydrogen peroxide system. Fenton’s reagent produces hydroxyl radical (0H') by the
following reaction:

2* + H202 -> 110' + Fe 3+ + 0H-

Fe

Hydroxyl radicals are powerful oxidizing agents. Fenton’s reagent has been used to

oxidize recalcitrant organic compounds such as PCBs (3), benzene, chlorobenzene (4),
diesel fuel (5), 2,4-D (6) and pyridine (7).

Silicone polymers have high chemical and biological stability. There are no

previously reported attempts to rnineralize polymeric compound such as

polydimethylsiloxane (PDMS; also refered to as silicone) or its monomer

dimethylsilenediol (DMSD) by Fenton’s reagent. If successful, the use of Fenton’s reagent

to decompose PDMS or DMSD could be applied to contaminated waters. There is also

the potential for utilizing Fenton’s reagent to remediate soil contaminated with PDMS or
DMSD. Iron is not generally a limiting factor for the reaction to occur in soil, however
hydrogen peroxide would need to be added. The objective of this project was to
determine whether and to what extent Fe/HzOz oxidation of PDMS and DMSD can

occur in water and soil.

Terminology

The term "silicone" is not used consistently. It is sometimes employed generically
to designate all monomeric and polymeric organosilicon compounds containing Si-C
bonds, and other times as a collective term for organosilicon polymers of all types (8). It
is also used strictlly to describe organosilicone polymers containing Si-O-Si bonds. This
usage will be adopted herein. Silicones occupy an intermediate position between organic
and inorganic compounds, and in particular between silicates and organic polymers. Few
chemical compounds are capable of so many applications as the Silicones. The 1991
domestic consumption of silicone fluids was estimated at 155-160 million pounds (9).
They are produced as oils, greases, resins and rubbers. They are used as engineering
materials in the construction of machines and other apparatus, as hydraulic fluids or heat
transfer media, as water-repellent agents in the textile and leather industry. The Silicones
are purportedly physiologically inert so they have also been used widely as auxiliaries in
medicine, pharmacy and cosmetics.

Polydimethylsiloxane (PDMS) is a part of a group of man-made organosilicone

compounds that range from low molecular weight volatile materials to high molecular

weight solids. The focus of this study is nonvolatile polymeric PDMS represented by the

generic formula:

 

 

CH3 CH3 CH3
I | l
CH3Si--O-- Si--O --SiCH3
I I I
CH3 CH3 " CH3
K J n

where n may vary from 10 to > 10 000.

PDMS has very low water solubility. The water solubility of trimethylsilyl capped
oliogomers, from the dimer to the pentamer, has been measured recently (10). Solubility
of the linear oligomers dropped from 930 ppb for the dimer to only 0.07 ppb for the linear
pentamer. Since the log of solubility was found to decrease linearly with increasing
molecular weight, the authors concluded that there is "an essential absence of any
ecologically significant water solubility for the conventional higher molecular weight

PDMS of commerce".

Environmental occurrence of PDMS

PDMS is used in a broad range of food, medical, personal care and household
applications (8). Its annual production is steadily increasing. These facts led to developing
concern over its potential adverse ecological impact due mainly to the following facts: (1)
PDMS fluids are essentially non reactive, except when catalyzed by strong acids and

bases; (2) PDMS based elastomers are uniquely resistant to weathering (i.e. hydrolysis

and oxidation); (3) no evidence has been found for biodegradation of these polymeric
fluids. In many of its applications, PDMS is disposed into wastewater treatment system
and has the potential to enter the environment as a component of effluents and sludges
from wastewater treatment plants.

The PDMS sorption constants (Kd) reported for a series of seven surface soils
were average 4,9002t1,250 11kg. The corresponding organic carbon matter normalized
sorption constants (Koc) were 590,000:510,000 ng (12). The small siloxanes and
water soluble silanols probably influence these values. Sorption constants for the large
polymers, which constitutes the majority of PDMS are probably much higher. A Koc of
590,000 means that PDMS is essentially immobile in soil.

A review of the environmental fate and ecological impact of organosilicon
materials suggested that they do not appear today to present any demonstrable ecological
threat (11). Studies of methylsiloxanes indicated that PDMS did not bioaccumulate or
exhibit significant environmental toxicity. Furthermore PDMS in sludge amended topsoil
present at concentrations of between 290 - 3500 mg/kg appeared to have no effect on (1)
seed germination or seedling survival, (2) cumulative loss of nitrate, (3) immature shoots,
mature shoots, mature grain/seeds or total biomass, and (4) number of soil
microorganisms including bacteria, fungi and actinomycetes. The quantity and timing of
sludge amendment, however, may affect loss of nitrate in leachates and soybean yields
(4).

Polyimethylsiloxane that enters wastewater treatment systems is expected to be
removed primarily by sorption to sludge solids given to its very low water solubility (l3,

l4). PDMS is expected to remain unchanged during wastewater treatment because of its

chemical and thermal stability (13). PDMS, not removed on sludge solids, will be present
in wastewater treatment effluent as a component of the suspended solids. It could be
deposited in sediments downstream of the facility. Most municipal sewage sludge today
is landfilled, incinerated, or applied to soil as fertilizer (Fig.1). Incineration converts
PDMS to water, C02 and Si. The predominant pathway of PDMS entry into the
environment is expected to be as a component of sludge added to the soil.

The earliest work of Bush et al. (16) indicated that PDMS fluids in many soils
undergo siloxane bond redistribution and hydrolysis, resulting in the formation of low
molecular weight cyclic and linear oligomers. Clays were shown to be the catalytic soil
component responsible for the above oligomerization and hydrolysis. In the absence of
clay minarals, PDMS is stable under mild environmental conditions (e.g., at ambient
temperature, pH 5-9, lack of strong oxidants). Once in contact with clay minerals, it does
degrade via the general pathway:

Me3Si0(SiMe20)nSiMe3(PDMS) —-> polymeric residue—9 mH0(SiMe20)xH + 2Me3Si0H
where Me is the methyl group; n>100; m<n-3x; with x=1-3, mainly 1.

The activity of all types of clays as catalysts for PDMS dagradation was shown to be
inversely related to the level of clay hydration. Furthermore, in the presence of
montrillonite attack was preferentially upon trimethylsilicone sites, leading primarily to
linear oligomerization, whereas kaolinite led predominantly to cyclic oligomers. These
results suggested the half-life of PDMS in soil could range from several minutes to weeks
or substantially longer. Relatively moist or wet soils exhibited significantly less catalytic

rearrangement activity.

Primary Treatment Secondary Treatment

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Influent Screws Grit ' Sedimentation Aeration Final
C hamba' Tank Tank Clarifier
Wastewater
Treatment

 

 

 

 

    
       
  

PDMS removed
by sorption to sludge
>94.0%

 

PDMS in
Effluent
<6.0%

 

 

Sludge
Disposal

V

 

 

 

 

   

Silanol volatizatio
biodegradation,

 
     

 

Landfill 34%
I   Other 6.4%

Surface Disposal 9.6%

    
    
     
 

  

Converted to
carbon dioxide
and silicone
dioxide

Incineration 16.0%
\

Figure 1. Overview of PDMS environmental mass flow as a result of wastewater
treatment practices. Fraction of sludge disposed by various techniques. Obtained from

EPA survey (15).

Similar results were shown by Xu (17) in a study aimed at investigating the ability of 12
different Ca—saturated clay minerals to catalyze the degradation of PDMS. Kaolinite,
beidellite and nontronite were most effective with 70-80% of the PDMS hydrolyzed to
DMSD within 24 hours. The variation in PDMS degradation rates among Ca—saturated
clays was thought to be due to the differences in the clay particle surface areas, surface
Bronsted and Lewis acidities and soil moisture. Higher surface area provided higher rates
and extents of PDMS degradation. In a related study, Xu (18) observed a 50% to 120
fold increase in hydrolysis when the exchangeable cations were changed from Na+ to Ca2+
or Al3+. Increasing the relative humidity from 32-100% significantly decreased
degradation of PDMS and led to the formation of volatile dimethyl cyclic compounds
such as oxtamethylcyclotetrasiloxane (D4). The suggested mechanism of degradation,
regardless of humidity, was via both random scission and end cutting. These studies
indicated that PDMS rearrangement products partition largely to the atmosphere.
Approximately 10 to 25 % of the original PDMS may partition to waterways as short
chain siloxanols. However, soil composition and hydration have a marked influence on
the amount and distribution of the rearrangement products.

Recent work by Lehmann et a1. (12, 19) has shown that once 14C-PDMS is
introduced into the soil environment, it can undergo extensive chemical hydrolysis to low
molecular weight water soluble products. Gas chromatography-mass spectrometry
identified the main product in all soils as dimethylsilane—l,1-diol (DMSD). Other small
silanols and cyclic siloxanes were either not detected or were present in only trace
amounts. No volatile 14C was captured by plugs of activated carbon on a polyester

support, and quantitative recovery of '4C showed no loss of unidentified volatiles. The

breakdown of PDMS was much more rapid under dry conditions, i.e. when soil moisture
content dropped below 5%. Carpenter et al. (14) demonstrated that within 4-6 weeks
more than half of the initially spiked silicone was converted to a water-soluble fraction
containing mostly DMSD.

Although sludge can be incorporated to 20-30 cm depths, PDMS will only
hydrolyze Extensively to DMSD in the top few cm, where soil drying is most likely.
PDMS hydrolysis at lower depths will occur more gradually because of higher soil
moisture (21). A similar transformation that occurs at a much slower rate has also been
observed in sediments (22).

An extensive monitoring program was conducted to determine the occurrence of
PDMS in environments impacted by waste disposal practices (23). Eight wastewater
treatment plants in North America were monitored to determine PDMS removal during
wastewater treatment. Surface waters, sediments, and sludge-amended soils were also
monitored for a complete assessment of the environmental fate of PDMS. PDMS sludge
concentrations reported were in the range of 290-5155 mg/kg and varied as a function of
influent concentration and sludge processing method. PDMS was found to be efficiently
removed from water during the wastewater treatment process (Figure l). Sediment levels
of <6mg/kg were measured near the outfall of the wastewater treatment plants. Measured
concentrations of sludge amended soils ranged from <0.4l to 10.4 mg/kg and were lower
than expected in most cases based on calculated PDMS loading via historical sludge
applications. The lower than expected PDMS concentration in sludge -amended soil
suggests degradation of PDMS in the soil environment to low molecular water extractable

fractions. According to all these studies, accumulation of PDMS is not likely to occur in

soils. Rather, the fate of Silicones may ultimately depend on the environmental fate of
DMSD.

Rapid degradation of a water soluble mixture of oligomeric dimethylsiloxanes by
sun light in the presence of environmentally relevant levels of nitrates was reported (24).
Silicate was demonstrated to be the final photodegradation product of these water soluble
siloxanes. When an aqueous emulsion of high molecular weight PDMS was" irradiated in
the presence of the nitrate ion a level of silicate corresponding to 1.7% of the original
quantity of PDMS was found after 5 weeks. This was explained to be due to the very low
solubility of PDMS and only a monomolecular surface layer of the polymer can react with
the hydroxyl radicals formed in the water. Thus droplet formation can be expected in the

relatively small irradiation cells resulting in a very small siloxane water interface area.

Environmental fate of DMSD
DMSD has the folloing structure:
CIH3
HO — Si — OH

CI.
The appearance of DMSD in soil represents hydrolysis of the PDMS to its monomeric
units. When the reaction of PDMS hydrolysis was investigated in variety of soils having
different pH, organic matter contents, textures and mineralogy, DMSD was the sole
reaction product (12). The fact that dimer, trimer, and larger siloxanols were not seen is

consistent with experiments showing that equilibrium in water favored the monomeric

form over the various oligomers (25).

Lehmann et. al. showed that DMSD could biodegrade and volatilize from soil
(26). Volatilization of’ ld’C-DMSD from soil was investigated in open tubes with weekly
remoistening (27). The loss of about 30% of the total l4C after seven weeks was assumed
to be due to volatilization. A portion of the 14C was also bound to the soil in a form
releasable by extraction with 0.1 M HCl. Some of the soil-bond 14C was released by
water after weekly remoisturing of the soil. These fractions are ithus not irreversibly
bound. DMSD was incubated in soils from permanent grass, crop field, deciduous
woodland, and pine woodland for 30 weeks at room temperature. After 30 weeks, most of
the remaining 14C in the soil was no longer water extractable but was acid and base
extractable. These fractions were not formed by microbial means. Incubations with 2%
plant litter showed reduced volatilization of the DMSD. In the presence of litter,
incubation with a microbial inhibitor resulted in less DMSD degradation. The cold
storage of soils almost completely stopped degradation, meaning that little transformation
of DMSD was expected during the colder months. DMSD is more likely to volatilize
from sandy than from clayey soils.

DMDS was found to be biodegradable in soils from residential gardens and sludge
disposal areas (28). The fungus F usarium oxysporum Schlechtendahl, and a bacterium, an
Arthrobacter species, were isolated from two different soils and both microorganisms
were able to cometabolize l4C-DMSD to 14C02. Active cultures of F. oxysporum
Schlecht. converted 19% of l4C-DMSD to 14CO2 after 240 days and cultures of the
Arthrobacter sp. converted more than 10% of it after 90 days.

After DMSD volatilizes into the atmosphere, it should undergo oxidation of its

methyl groups by reaction with sunlight-induced hydroxyl radicals in the air (in a similar

10

fashion to other volatile siloxanes) (29, 30). If the compound is washed out of the air in
rainfall, it may be oxidized in water by photolytic demethylation (2, 31), or it could be
microbially oxidized in soil (27). In either event, expected degradation products of
DMSD are CO2 and silicic acid. The latter will precipitate to form amorphous soil

minerals (32).

F enton chemistry

In April 1876, Henry John Fenton (then an undergraduate at Chist’s College,
Cambridge, England) wrote a letter to the editor of Chemical News (33) reporting a
reaction between ferrous sulfate, hydrogen peroxide (H202) and tartaric acid. The purpose
of the letter was to describe a color test for tartaric acid. Later it became evident that this
mixture, called Fenton reagent, was an effective oxidation agent for various organic
compounds. Fenton’s first report (33) and his full paper (34) demonstrated three key
features of what we now term "Fenton Chemistry": (‘1) an oxidant (2) a reduced form of a
heavy metal (in this case iron); and (3) a higher oxidation state of iron involved as an
intermediate. The reaction is believed to proceed as follows:

Fe (11) + H202 —> Fe (HI) + 'OH + OH'0 (1)

A few years after Fenton’s death in 1929, Haber, at al. (35) proposed the
involvement of hydroxyl radical in the iron-catalyzed decomposition of hydrogen
peroxide: They proposed the following scheme:

'OH + H202 —) HOO'2 + H2 (2)
H00' + H202 —> 02 +‘0H + H2 (3)

HOO‘ 4:» H+ + O' 2 (4)

ll

often described as the "Haber-Weiss" reactions. Equation (3) is thus the potential source
of 'OH from superoxide/peroxide in the absence of metals. The superoxide radical HO‘ 2
is known to have lower oxidation capability compared to the hydroxyl radical.

Barb et. al. (36) later studied the decomposition of hydrogen peroxide catalyzed
by ferrous and ferric ions and modified the original Haber-Weiss mechanism in the

following way:

Initiation: Fe 3" + H202 —-> Fe2+ + HOO' + H+ (5)

Propagation: Fe2+ + H202 —> Fe 3" +‘0H + OH“ (6)
‘OH + H202 —> H20 + H00' (7)
HOO‘ + Fe3+ —) Fe2+ + H+ + 02 (8)

Termination: HOO' + Fe2+ —2 H00' + Fe3+ (9)
'OH + Fe2+ —> no + Fe3+ (10)

An important finding was that the oxidation effects of Fenton reagent strongly depended
on the concentration ratio of H202 and Fe”. If an organic compound is present in the
system, the reaction includes two more steps:

RH+‘OH—>H2O+R (11)

R' + Fe3+ —-) Fe2+ + products (12)

The parallel application of H202 photolysis and Fenton reaction has lead to the
substantial improvement in efficiency of the pollution degradation. In the Fenton reaction,
Fe 3+ ions are accumulated in the system and, after the Fe2+ ions are exhausted, the
reaction terminates. To avoid this problem, the photochemical regeneration of ferrous

ions by photoreduction of ferric ions was proposed. It has been known for a long time that

12

the photochemical irradiation of ferric species leads to their reduction (37). The following
reaction proceeds in water solution:

Fe” + H2O + hv—> Fe 2* + HO‘ + H+
It was further ascertained that Fe (III) ions accelerate the photodecomposition of hydrogen
peroxide, thus catalyzing more efficiently the photochemical oxidation of different

compounds.

Application of the Fenton reaction

Recently, considerable interest has developed in utilizing of the Fenton reaction
for environmental remediation. Numerous transformation reactions take place in the
environment. In natural waters exposed to solar radiation, reactive intermediates are
formed which than take part in photo-oxidation reactions. They are known to be primary
sources of photochemical self cleaning of the atmosphere and aquatic environment (38).
The extent to which the Fenton reaction occurs in natural waters depends on the
concentrations of Fez“: Fe“, Cu2+ and H202. This chemistry has recently been applied to
the destruction of environmental pollutants (39, 2).
The aim of these reactions is to attain complete mineralization of the contaminants to
innocuous end-products such as CO2 and water. The extent of mineralization via Fenton’s
reaction can be measured by the final value of COD (chemical oxygen demand) or TOC
(total organic carbon). Systems generating hydroxyl radicals are among the most efficient

ones used for destruction of pollutants due to the high reactivity of the radical species.

13

Utilization of the F enton reaction in water treatment

A growing number of published papers in this field demonstrate the suitability of
the application of the Fenton reaction in water treatment. The use of the Fenton reagent is
advantageous over some other treatment procedures mainly because H202 decomposes to
harmless substances. Also only catalytic amounts of Fe2+ are used in the reaction and the
equipment needed for this treatment process is technically simple. Another advantage of
this process is that reactive hydroxyl radicals generated in the Fenton reaction can
mineralize recalcitrant pollutants. It was found (40) that the reaction yield strongly
depends on pH. At low pH, Fe(OH)+ intermediate contributed to additional formation of
hydroxyl radicals. The reaction rate drops off sharply above and below the optimum pH
of about 2.8. At higher pH the loss of reactivity is due to precipitation of Fe (111) as
hydrous oxyhydroxide, Fe203.n H2O

The Fenton reaction is most often used either in the pretreatment of waters with a
high value of COD prior to biological treatment or in the mineralization of toxic or
difficult to degrade pollutants. Heavily polluted and biologically harmful wastewater
from pharmaceutical industry was successfully pretreated in this fashion for biological
treatment. The total effect of the pretreatment was 92-95% decomposition, in contrast to
the conventional processes used (36). The Fenton reagent is frequently used for the
degradation of chlorinated organic pollutants. Koyoma et al. investigated the degradation
of 10 chlorinated aromatic compounds (39). The Fenton reaction was applied in the first
stage of degradation process followed by methanogenic fermentation using anaerobically
digested sludges. The results indicated that the chlorinated aromatic compounds were

degraded completely after Fenton reaction and some of them were even mineralized.

l4

Oxalic and formic acids were isolated as the main semiproducts of the oxidation.
Application of the Fenton reaction in the degradation of nitrobenzene and nitrophenols
was also investigated (41). The Fenton reaction was much more efficient in the
degradation of these compounds than photolysis in the presence of H202.
2,4-dinitrotoluene was completely mineralized within 5 hours at molar ratio
2,4-dinitrotoluenezH202 : Fe2+ of 1:20:25 (2).

The efficacy of the Fenton reaction is substantially improved in the photo Fenton
reaction. This consists of the parallel application of classical Fenton reaction and
photolysis. Presently the photo Fenton reaction is used more often in water treatment
processes. Sun and Pignatello (3) compared the efficiencies of the Fenton and photo
Fenton reactions in the degradation of 2,4-dichloroacetic acid (2,4-D) and
2,4,5-trichloroacetic acid (2,4,5-T). They showed that the classical Fenton reaction
catalyzed by Fe2+ was more efficient than the reaction catalyzed by Fe3+ ions. Application
of radiation at wavelengths above 300 nm significantly increased the rate of
mineralization. The consumption of H202 was reduced in this process to 5 mol peroxide
per 1 mol herbicide. Photo-Fenton reaction was also used for mineralization of 4-
chlorophenol in water (42). The authors suggested that the applied radiation (320) nm
substantially accelerated the degradation of 4—chlorophenol in the presence of H202 and
Fe2+. It was stressed that an advantage of photo-Fenton reaction was the small amount of
iron used as a catalyst and the low loading of the treated water. Pignatello et al. (43)
reported successful degradation of the commercial PCB mixture Arochlor 1242 in
acidified water at 66°C. Up to 88% PCB removal and 85% dechlorination were obtained.

Perfluorinated surfactants added to increase PCB solubility or rates of dissolution had a

detrimental effect on degradation. Minimal reaction was observed at room temperature
owning to the low water solubility of PCBs and lower rate of generation of hydroxyl
radicals.

Fenton’s reaction has been shown to be temperature dependent. Two studies
performed to follow the effect of temperature on the degradation of methyl parathion (44)
and 2,4-dinitrotoluene (2) showed that the loss of the compounds increased from about
20% at 10°C to about 90% at 35°C, with no further improvement up to 60°C.

Recently, Rabek et al. reported a project aimed to implement Fenton and
photo-Fenton reaction for destroying water-soluble polymers like poly (acrylic acid) and
poly (ethylene dioxide) (45). The results showed that the process is completed in 15 to 30
min. In the presence of dissolved oxygen a chain scission reaction occurred with a

concurrent decrease in solution viscosity.

Utilization of the Fenton reaction in soil treatment

Contaminants in soils are often sorbed or present as nonaqueous phase liquids.
This can have significant impact on the effectiveness of Fenton’s process chemistry.
Watts et a1. (46) first reported the use of a Fenton-like reaction for soil remediation and
showed that pentachlorophenol in silica sand slurry was destroyed at pH 2-3. They
confirmed that when high concentrations of H202 were used for soil remediation, soluble
iron concentrations decline rapidly over first 30 minutes of the reaction with subsequent
formation of an amorphous floc over the following three hours. In a separate study
Fenton’s reagent was used to oxidize pentachlorophenol, trifluralin, hexadecane, and

dieldrin in soils, and it was found that iron oxyhydroxides effectively catalyzed

l6

Fenton-like reactions (47). The most effective oxidation of soluble pentachlorophenol in
silica sand was catalyzed by naturally occurring iron oxyhydroxides, however sorbed
pentachlorophenol was not oxidized with goethite as the only Fenton’s catalyst (5). These
results suggest that naturally occurring iron oxyhydroxides can serve as effective Fenton
catalysts. A possible mechanism for mineral catalyzed Fenton’s reaction has been
proposed by Kitajima et al. (48):
H2O2+S -> S++ 'OH + OH',

where S is the mineral surface and S+ is an oxidized region on the mineral. These
heterogeneous mineral catalysts are effective because they provide consistent rates of
hydroxyl radical generation when high concentration of H202 are used.

In most of the soil remediation studies using Fenton’s reaction, the soil pH was
adjusted to about 3.0 with 0.1 N H2804 to reach the effective pH range (1). Acidification
of the soil to the proper pH however is cumbersome and may preclude or severely restrict
implementation of the Fenton’s reaction for in situ treatment. Moreover acidification and
subsequent neutralization will add two additional steps to any scheme. In order to avoid
this Sun and Pignatello (40) tested a number of structurally diverse organic and inorganic
polydentate chelators for their ability to solubilize Fe (III) at pH 6.0 and catalyze
oxidation of 2,4-D in aerated aqueous solutions. Most of the compounds tested that
contained organic ligands, formed soluble Fe (III) and were capable of oxidizing 2,4-D in
periods from minutes to less than 5 hours. All active organic ligands were themselves
oxidized, but in many cases the resulting iron complexes remained soluble and reactive.

From a practical perspective, degradation of the ligands may be viewed as advantageous

17

because provided iron remains soluble and catalytically active long enough to carry out
degradation.

Iron complexes have also been shown to catalyze the Fenton-like oxidation of
pesticides. The primary advantage of the iron complexes is the potential for effective
generation of hydroxyl radical at near-neutral pH. Pignatello and Day (44) screened
several iron (III) complexes for their potential to promote the oxidation of pesticides by
Fenton-like reactions. Nitrilotriacetic acid, tetraxydroxy-l,4-quinone and
hexaketocyclohexane were among the most effective ligands evaluated. Watts et al. (47)
treated diesel-contaminated soils with H202 and six iron compounds and naturally
occurring minerals. Iron (H) compounds were less effective catalysts than iron (III)
species owing to the initial loss of H202 in oxidizing iron (H) to iron (11]). The most
effective iron salt catalysts were:

C104‘ = NO3'> 30,25 PO41

Watts et al. (49) applied Fenton chemistry to decompose pentachlorophenol in
soils. They used H202 concentrations ranging from 3 to 300 mM and iron (11) sulfate as a
catalyst rather than natural occun'ing iron oxyhydroxides. Oxidation of sorbed to soil
pentachlorophenol was effective at H202 concentrations as low as 30 mM if catalysts
were presented as soluble iron rather than natural iron oxyhydroxides. Because of the
high relative concentration of the iron mineral catalysts in a slurry with a low volume
ratio of liquid to soil, H202 decomposition was rapid (i.e., < 8 h.). Therefore, the use of
soluble iron to catalyze the Fenton-like oxidation of sorbed contaminants may provide

more economical and effective process conditions.

18

REFERENCES

1. Spencer, C.J., P.C. Staton, R.J. Watts, A central composite rotatable analysis for the
catalized hydrogen peroxide remediation of diesel contaminated soils, J .Air Waste
Manage.Assoc., 46: 1067, 1996

2. Mohanty, N.R., I.W.Weil, Oxidation of 2,4-Dinitrotoluene using F enton ’s reagent.
Reaction mechnisms and their practical applications., Haz.Waste Haz. Mater., 10:171,
1993

3. Sun, Y., 1.]. Pignatello, Photochemical reactions involved in the total mineralization of
2,4-D by Fe3+/l1202/UI/, Environ.Sci.Technol., 27:304, 1993

4. Tolle, D.A., C.L. Frye, R.G. Lehmann, T.C. Zwick, Ecological Effects of PDMS
augmented sludge amended to agricultural microcosms, Sci. Total Environ., 1622193-
207, 1995

5. Watts, R.J.,M.D.Udel, R.M.Monsen, Oxidation of pentachlorophenol catalyzed by
H202 and naturally occurring iron oxyhydroxides, Water Environ.Res., 652893, 1993

6. Barberini, M., C. Minero, E. Pelizzetti, E. Borgarello, N. Serpone, Chemical
degradation of chlorophenols by F enton ’s reagent., Chmosphere, 1612225, 1987

7. Watts, R.J., S.E.Dilly, Evaluation of iron catalysts for the F enton-like remediation of
diesel contaminated soil, J.Haz.Mater., 511209, 1996

8. Noll, W., Chemistry and Technology of Silicones, Academic Press Inc., New York, NY
10003, 1968

9. Chemical Economics Handbook, SRI International, 1990

19

10. Varaprath, S.,C.L.Frye, J.Hamelink, Aqueous solubility of permethylsoloxanes
(silicones), Environ.Toxicol.Chem., 8: 1263, 1996

ll. Frye, C.L., Environmental fate and ecological impact of organosilicon materials: a
review, Sci. Total Environ, 73: 17-22, 1988

12. Lehmann, R.G., S.Varaprath, .J.L.Arndt, Degradation of silicone polymers in a
variety of soils, Environ.Toxicol.Chem., 8: 1299, 1995

13. Watts, R.S., S.King, C.S.Haling, L.Gearhart, B.Vigon, Water Res., 29:2405, 1995

14. Lentz, C., Silicones in the environment, Ind.Res.Dev., 22: 139, 1980

15. Standard for use and disposal of sludge, Fed.Regist., 58:9257, 1993

16. Buch, R.R., D.N.Ingerbrigsten, Rearragement of Poly(dimethylsiloxane) fluids on
soil, Environ.Sci.Technol., 6:676, 1979

17. Xu, Shihe, R.G. Lehmann, JR. Miller, G. Chandra, Degradation of PDMS (Silicones)
as influenced by clay minerals, Environ. Sci. Technol., 32: 1 199-1206, 1998

18. Xu, Shihe, Hydrolysis of PDMS (silicone) as catalyzed by clay minerals at two
diflerent moistures, Environ. Sci. Technol., Submitted for publication

19. Lehmann, R.G., S.Varaprath, C.L. Frye, Degradation of silicone polymers in soil,
Environ.Toxicol.Chem., 7: 1061, 1994

20. Carpenter, J.C., J.A.Cella, S.B.Dom, Study of degradation of PDMS on soil,
Environ.Sci. Technol., 291864, 1995

21. Lehmann, R.G., C.L.Frye, D.A. Tolle, T.C.Zwick, Fate of sludge-applied silicones in

agricultural soil microcosms, Water,Air Soil Poll., 83: 1, 1995

20

22. Dom, S.B., S.Frame, Development of high-performance liquid chromatographic-
inductively coupled plasma method for speciation and quantification of silicones: from
silanols to polysiloxanes, Analyst, 119:1687, 1994

23. Fendinger, N.J., D.C.McAvoy, W.S.Eckhoff, B.B.Price, Environmental occurance of
PDMS, Environ.Sci.Technol., 31: 1555,1997

24. Buch, R.R., T.H.Lane, R.B.Annelin, C.L.Frye, Photolytic oxidative demethylation of
aqueous dimethylsiloxanols, Environ.Toxicol.Chem., 6:676, 1984

25.Spivac, 1., SB. Dom, Hydrolysis of oligodimethylsiloxane-a (1)-dials and the position
of hydrolytic equlibrium, Environ. Sci. Technol., 2822345, 1994

26. Lehmann, R.G., C.L. Frye, Fate of silicone dagradation products in soil, Environ.
Toxicol. Chem., 13:1753, 1994

27. Lehmann, R.G., J.R.Miller, Volatilization and sorption of DMSD in soil, Environ.
Toxicol. Chem., 9: 1455, 1996

28. Sabourin, C.L.,J.C. Carpenter, T.K. Leib, J.L. Spivack, Biodegradation of
dimethylsilanediol in soils, Appl. Environ. Microbiol., 62:4352-4360, 1996

29. Atkinson, R., Kinetics of the gas-phase reactions of a series of organosilicon

compounds with . OH and N0. 3 radicals and 03 at 295K, Environ.Sci.Technol., 252863,
1991

30. Sommerlade, R., H.Parlar, D.Wrobel, P.Kochs, Product analysis and kinetics of the
gas-phase reactions of selected organosilicon compounds with OH radicals using a smog
hamber-mass spectrometer system, Environ.Sci.Technol., 27:2435, 1993

31. Anderson, C.R., Studies of the oxidative photoinduced degradation of silicones in

aquatic environment, Chemosphere, 16:2567, 1987

21

32. Lindsay, W.L., Chemical Equilibrium in Soils, John Wiley, New York, NY, USA,
1979

33. Fenton, H.J.H., On a new reaction of tartaric acid, Chem.News, 28:39, 1876

34. Fenton, H.J.H., Oxidation of tartaric acid in presence of iron, Chem.Soc., 65:899,
1894

35. Haber, F., J. Weiss, Uber die katalyse deshydroperoxide, Naturwissenschaften,
202948, 1932

36. Barb, W.G., J.H. Baxendale, P.George, KR. Hargrave, Reactions of ferrous and
ferric ions with hydrogen peroxide, Trans.Faraday Soc., 472462, 1951

37. Prousek, J., F enton reaction after a century, Chem.Lisy, 89:11, 1995

38. Prousek, J., Advanced oxidation processes for water treatment. Photochemical
processes, Chem. Lisy, 872165, 1993

39. Koyama, 0., Y.Kamagata, K.Nakamura, Degradation of chlorinated aromatics by
F enton oxidation and methanogenic digester sludge, Water Res., 28:895, 1994

40. Sun, Y., J.J.Pignatello, Chemical treatment of pesticide waste. Evaluation of F e(III )
chelate complexes catalytic H202 oxidation of 2,4-D at curcumneutral pH, J.Agric.Food
Chem., 40:322, 1992

41. Lipczinska-Kochany, E., G.Sprah, S.Harms, Influence of some groundwater and
surface waters constituents on the degradation of 4-chlorophenol by the F enton reaction,
Chemosphere, 3029, 1995

42. Ruppert, G., R.Bauer, G.Heisler, The photo-Fenton reaction - an eflective
photochemical wastewater treatment process, J.Photochem.Photobiol.A. Chem., 73:75,

1993

22

43. Pignatello, J.J., G.Chapa, Degradation of PCBs by ferric ion, H202 and UV light,
Environ.Toxicol.Chem., 13:423. 1994

44. Pignatello, J.J., Day, Mineralization of methyl parathion insecticide in soil by
hydrogen peroxide activated with iron(IIl)-NTA or HRIDA complexes,Haz. Waste Haz.
Mater., 13:237-244, 1996

45. Linden, L.A., J.F.Rabek, H.Kaszmarek, A.Kaminska, M.Scoponi, Photooxidative

degradation of polymers by 110‘ and H02. radicals generated during the photolysis of
H202, FeCl3 and F enton reaction, Coord.Chem.Rev., 125: 195, 1993

46. Watts, R.J., M.D.Udel, P.A.Rauch, S.W.Leung, Treatment of Pentachlorophenol-
contaminated soils using Fenton ’s reagent, Haz.Waste Haz.Mater., 72335, 1990

47. Tyre, B.W., R.J.Watts, G.C.Miller, Treatment of four biorefractory contaminants in
soils using catalyzed hydrogen peroxide, J.Environ.Qual., 202832, 1991

48. Kitajima, N., S. Fukuzumi, Y.Ono, Formation of superoxide ion during the
decomposition of hydrogen peroxide on supporetd metal oxides, J.Phys.Chem., 82: 1505,
1978

49. Watts, R.J., S.Kong, M.Pippre, Treatment of contaminated soils using catalyzed

H202, Chemical Oxidation, 39:33, 1994

23

CHAPTER I

INTRODUCTION

Dimethylsilanediol (DMSD) is the major product of the polydimethylsiloxane
(PDMS) hydrolysis in soil (1,2). PDMS is introduced in the environment as component of
municipal sewage sludge applied to soil. Previous research has demonstrated that PDMS
may undergo extensive degradation in dry soil (3). The clay constituents of the soil were
found to promote this process but degradation was found to be markedly inhibited by
moisture. In soil, PDMS is converted to a water soluble fraction containing mostly
dimethylsilanediol (DMSD). DMSD has the structure 0H-Si-(CH3)2-OH. Once formed,
DMSD may remain bound to soil particles, undergo complete biodegradation to CO2 or
volatilize in the atmosphere and be oxidized photochemically.

Fenton reagent has proved to be an expensive and powerful oxidant that has been
shown to oxidize a wide variety of organic compounds (4, 5). In the Fenton’s reaction
both Fe(H) and Fe(III) react with hydrogen peroxide (H202) to produce hydroxyl radicals
('OH). The DH radicals are capable of oxidizing a range of contaminants because they
react with most organic compounds in solution at near diffusion controlled rates (6). The
use of Fenton’s reagent for oxidizing biorefractory organic compounds in aqueous
solution and contaminated soil has been studied extensively (7, 8).

The purpose of this research was to evaluate the potential for DMSD degradation

by Fenton reagent in water and soils and to optimize the reaction conditions based on ( 1)

24

the concentration and oxidation state of the iron catalyst, (2) the chemical form of the iron

catalysts, (3) H202 concentration and (4) the sluny volume.

MATERIALS AND METHODS

Materials:

Dimethylsilanediol (DMSD), ”C-DMSD, polydimethylsiloxane (PDMS) and l“c-
PDMS with 99% radiochemical purity were provided by the Dow Corning Co. (Midland,
MI). Hydrogen peroxide and ferric nitrate were purchased from J.T.Baker (Phillipsburg,
NJ). Ferrous chloride, ferric chloride, ferrous peroxide and ferric peroxide were obtained
from Fisher Scientific (Fail Lawn, NJ). All other chemicals used (NaOH, HCl, methanol)
were analytical grade.

Samples of the Capac sandy loam A-horizon soil were collected in East Lansing, MI.
Particle size distribution was determined by the pipette method (9). Organic carbon
content was measured by combustion at 900°C with C02 trapped in KOH (10). Cation
exchange capacity (CEC) was determined by sodium acetate saturation at pH 8.2 (11).
Crystalline and amorphous iron and manganese concentrations were quantified by citrate-

dithionate-extractions (12). The soil characteristics are given in Table 1.

Table I-l. Characteristics of the Capac A soil

 

 

Characteristic Value Characteristic Value
Organic carbon (%) 3.28 Crystalline Iron (mg/kg) 12404
Sand (%) 54.6 Crystalline Manganese (mg/kg) 239
Silt (%) 24 Amorphous Iron (mg/kg) 8481
Clay (%) 21.4 Amorphous Manganese (mg/kg) 136
pH 6.85

CEC ( meq/lOOg) 43.0

 

Methods:
DMSD - \Mater experiments:

A slightly modified version of the method described by Ronen et al. (13) was used to
evaluate the decomposition of DMSD in water by Fenton reaction. Reactions were run in
60 ml serum bottles containing a reactant volume of 10 ml. All reactant were prepared in
deionized (DI) water. A stock solution of 5 mM DMSD/ ‘4 C- DMSD were prepared in
advance, and introduced into the reaction vessel. The amount of 14 C- DMSD for each
bottle was about 5 nCi. Stock solutions of ferrous and ferric salts (20 mM) were made in
fresh in DI water acidified with HCl, and introduced into the reaction vessel. The iron (H)
chloride and iron (III) chloride concentrations were varied from 0.1 to 10 mM and the
H202 concentrations ranged from 5.5 to 880.0 mM. The pH of the reaction solution
dropped below 3.0 after addition of H202. The influence of the initial DMSD
concentrations on reaction kinetics and efficiency were evaluated between 0.1 and 3.0

mM DMSD. Other iron catalysts such as iron (H) sulfate, iron (III) sulfate, iron (H)

26

perchlorate, iron (1H) perchlorate and iron (HI) nitrate were used to catalyze the reaction
with final concentration in the reaction media of 2 mM. Bottles were closed with rubber
stoppers fitted with center well (Kontes, Morton Grove, IL) containing a strip of 5-x-1 cm
filter paper and 0.2 ml 1N NaOH used to trap l4C02. Reactions were performed in the
dark at room temperature with shaking at 100 rpm. At specific intervals the filter paper
was removed and placed in a vial with 10 ml scintillation cocktail. Vials were left
overnight in the dark and then radioactivity was measured by liquid scintillation counting
(LSC). An aliquot (1 ml) of the reaction solution was added to a separate scintillation
vials with 10 ml cocktail and the radioactivity was counted using LSC. Each treatment
was performed in duplicate, and all results are expressed as a mean value i standard

error. The range of the standard errors was 3.0-5.5%.

DMSD-Soil experiments:

Air dry (2.5 g) soil was weighed in 60 ml serum bottles and amended with 5mM
DMSD/l4C-DMSD stock solution to yield initial concentrations of DMSD between 1-200
mg/kg. Iron concentrations ranged between 50 and 150 mM and hydrogen peroxide
concentrations varied from 0.1 to 3.0 M; H202 concentrations greater than 3.0 mM were
not used because the slurries produced a violent reaction. The total volume of the liquid
was 2.5 ml. The contents of the serum bottles were mixed with a vortex mixer and then
shaken at 200 rpm at dark for various reaction times at room temperature. The same
procedure as that used in the water experiment was followed for sampling thetraps. After
completion of the reactions the soil was transferred to centrifuge tubes. Tubes were

shaken for 6 h with 5 ml DI water and then centrifuged at 8500 rpm for 20 minutes. The

27

water was removed and duplicate l-ml aliquots were sampled for 14C analysis. Tubes
were than shaken overnight with 5 ml 0.1 M HCl, centrifuged, and lml of the liquid was
sampled for 14C as above. To investigate the most efficient form of the iron catalyst
different iron species were added to the soil at a concentration of 100 mM/kg. The forms
of iron used were Fe(H) chloride, Fe(HI) chloride, Fe(H) perchlorate, Fe(Hl) perchlorate,
Fe(H) sulfate, Fe(HI) sulfate, Fe(HI) nitrate. The reactions were allowed to proceed for
24h. It was found that after 5 hours, reactions were completed because H202 was
consumed. In order to drive the reactions further additional H202 was added to the
slurries after the 5th hour at the same initial concentration.

The effect of the slurry volume was investigated using the most effective catalysts
and its optimum concentration. The catalyst Fe(NO3)3 at a fixed concentration of 100 mM
was used with H202 concentrations of 0.625 M and liquid volume from 1.5 to 5 ml per
2.5 g of soil.

All experiments were initiated by addition of H202 to the vials. Control experiments
using DMSD alone, combinations of DMSD and H202, and combinations of DMSD and
iron species were performed in parallel. Each treatment was performed in duplicate, and
all results are expressed as a mean value i standard error. The range of the standard errors

was 2.8-5.0%.

28

RESULTS AND DISSCUSION

Exposure of DMSD to Fenton reagent in water resulted in the mineralization of
the chemical. During incubation for 24 hours 75% of DMSD was mineralized in water
and 48% in soil as determined by the formation of '4 CO2. Control systems with DI.
water, H202 and iron (H) or iron (IH) chloride alone conducted in parallel were not
capable of mineralizing the DMSD. After completeion of the reactions the radioactivity
remaining in the water was less than 0.5% for the Fenton treated samples; volatilization
loss of DMSD was about 25%. For the control experiments the recovery ranged between
90.7 and 100.6%. The initial concentration of DMSD, H202 and Fe (H)/Fe (H1) ions as

well as the iron catalyst source influenced the mineralization of DMSD.

Mineralization of DMSD by Fenton’s reagent in water

The transformation of DMSD by Fenton’s reagent was practically independent of
initial pH of the reaction solution. The pH dropped to 2.8-3.0 within seconds on addition
of H202, presumably due to release of protons from hydrolysis of the Fe (H1) products.
These observations are consistent with previous reports regarding the effect of pH on
Fenton oxidation of organic pollutants (14) in unbuffered water solutions. They
immediately became acidic upon addition of the reagents and thus post treatment
neutralization may be needed. Our system was also unbuffered because the presence of
0.1-0.02 M borate, phosphate and universal buffer drastically reduced the mineralization
of the DMSD (up to 20%) or completely stopped it. Similar results were demonstrated by

Lipczinska—Kochany et al. (15) in a study of the effect of several anions in concentrations

29

close to those found in groundwater on Fenton degradation of 4-chlorophenol. The
following order of anions was shown to inhibit the reaction:

(3104' ~ NO3’ < so} < Cl' << HPO4' < HCO3'
The observed retardation was explained as the combined effect of the anions on the rate
of oxidation of Fe (H) anions as well as their role in the scavenging of hydroxyl radicals.
It is well known that phosphate and bicarbonate are efficient scavengers of the hydroxyl
radical ( 16).

A major aim of this study was to optimize the Fenton reaction conditions for
DMSD mineralization in water as well as to compare the efficacy of Fe (H) and Fe (IH)
ions as catalysts. The chemical mineralization of DMSD was affected by the initial
concentrations of the iron catalyst and its oxidation state. Figure I-l shows the effect of
Fe (H) and Fe (IH) concentrations on decomposition of DMSD. In water the optimal
concentration of Fe (H) chloride needed to reach maximum conversion of DMSD to CO2
was reached at 2 mM (Figure I-l), was constant up to 4 mM, and slightly diminished at 8
and 10 mM. The optimal concentration of Fe (HI) chloride was between 1-4 mM (Figure
l-2). At the optimum Fe (H) concentrations, about 70 % DMSD removal was obtained
after 1 to 10 hours treatment. For Fe (H1) the time needed for the reaction to be completed
was 1 to 5 hours. Within the optimal concentration ranges the DMSD decomposition rate
during the first 5 hours increased with increasing Fe (H)/Fe (IH) concentrations, perhaps

due to the increased hydroxyl radical formation rate. Perhaps at higher than optimal iron

concentrations the initial formation rate of .OH radicals originating from the
decomposition of H202 was so high that they were consumed by competing side reactions

before being utilized effectively for DMDS mineralization.

30

 

 

 

 

 

X0 UK

 

 

 

 

 

 

’5
0)
'U
'U
(U
*
O
o\°
or
O . +
m 80 _____-—— / __ W ,,__ W
30 'I
70 T A_—-————"" ‘7’: +9:
* *
a
FeC I3
/+

 

 

 

 

 

' l
o 5 1O 15 20 25
Reaction tim e (hour)

Figure I-l. Mineralization of DMSD in water by Fenton’s reagent with different
iron catalysts concentrations (lmM DMSD, 440 mM H202). Treatments consisted of (I)
DMSD only, (0) H202 only, (A ) 0.1 mM iron only,( v) 2 mM iron only, (9) 10 mM
iron only, (+) 0.1 mM iron, (X) 1 mM iron, (*) 2 mM iron, (A) 4 mM iron, (0) 8 mM

iron, ([21) 10 mM iron.

31

 

 

 

 

 

 

 

 

14C as CO2 ( % of added)

 

 

 

 

 

1.5-I
1.0- /’:;//
05 ‘ —//;
' J /. .r
0.0 . .4 T g l I 41 I I I
0 5 10 15 20 25

Reaction time ( hour)

Figure I-2. Mineralization of DMSD in water by Fenton’s reagent with different
iron catalysts ( lmM DMSD, 2mM Fe(H)/Fe(HD, 440 mM H202). Treatments consisted of
(a) Fenton’s reagent samples, (b) controls. (I) FeCl2, (O) FeCl3, (A ) FeSOa,

(v ) Fe2(SO4)3, (O) Fe(C104)2, (+) Fe(ClO4)3, ( X) Fe(NO3)3, (CI) DMSD only, (0) H202

only.

32

Therefore, one possible reason for the existence of an optimum Fe (H)/Fe (H1)

dosage was the need to attain a proper .OH radical formation rate.

Another potential factor was the formation of iron hydroxide precipitate. The
ferric ions exist mainly in the form of soluble Fe(OH)2Jr at pH 3.0. Iron hydroxide
precipitates would form at concentrations slightly higher than 1 mM, and iron in this form
could not be effectively utilized as a catalyst (17). At the higher than optimal iron
concentrations the reaction solution had persistent deep red color formed after several
hours, which is characteristic of soluble polynuclear ferric oxyhydroxides (18). The faster
reaction kinetics observed for the Fe (HI) catalyzed reaction may be due to the additional
time needed for Fe (H) to be oxidized to Fe (H1). The Fenton reaction nominally between
Fe (H) and H202 under conditions of excess peroxide (as in our case) may be actually
predominantly the reaction between Fe (HI) and H202 since Fe (H) is rapidly oxidized.

The degradation reactions were sensitive to the anion in solution originating from
background electrolyte and Fe (H)/Fe (IH) or H+ counterions. Seven iron species were
evaluated for their ability to catalyze DMSD degradation. The data for mineralization
show that Fenton reaction catalyzed by iron (IH) was completed in shorter times
compared to the reaction catalyzed by Fe (H) species. All of the iron salts tested resulted
in about 70% mineralization of DMSD. However, differences were found in the initial
rates (Table 1-2). The initial reaction rates for Fe(H) catalysts were detrmined by using the
initial experimental portion of the experimental kinetic curves and compared to the initial

rates calculated using the first order kinetic model:

33

ln{Cmax/[Cmax-C(t)]} = kt,
where Cmax is the final CO2 percentage, C(t) is the CO2 persentage at
time t, and k is the rate constant.
The fitting procedure was carried out by the computer program “SigmaPlot” ( version
4.0 for Windows, by SPSS Inc., 1997 ). The regression parameter R2 was in the range
99.86 to 99.99%. First order kinetic model did not describe properly the data for Fe(Hl)

catalyzed reactions.

Table I-2. Kinetic parameters for DMSD degradation in water with different iron

 

 

 

 

 

 

 

 

 

catalysts.
Iron Initial rates Tl/2, hour Initial rates T1/2, hour
Catalyst mmol.L'l .sec"l (experimental) mmol.L’ ‘ .sec"l (calculated)
(experimental) (calculated)

FeCl2 14.3 3.0 15.8 3.0

FeCl3 66.6 0.5 -

FeSOa 32.0 2.4 26.9 2.0
Fe2(804)3 43.1 0.7 -

Fe(C104)2 7.9 . 4.9 9.3 5.0
Fe(C104)3 59.6 0.5 -

Fe(NO3)3 45.7 0.4 -

 

 

 

 

 

 

34

The most effective iron catalysts were Fe(C104)3 and FeCl3 followed by the chloride and
sulfate salts (Figure I-2 (a) and Figure I-2 (b). The classical Fenton reagent FeCl2 /H2O2
produced the slowest initial rate of mineralization of DMSD, followed by FeSO4. These
results indicate that to achieve mineralization there is little or no advantage starting with
Fe (H) under conditions with excess peroxide, as used here.

The concentration ratio of H202 to DMSD also influenced the effectiveness of
DMSD mineralization. The effect of H202 on the mineralization of DMSD in water is
evident at lower concentrations. The degree of DMSD mineralization in the presence of
FeCl3 increased steadily when the H202 concentration increased between 5.5 and 110 mM
(Figure I—3). However the mineralization rate did not increase when H202 concentration
was further between 110 and 880 mM. In the presence of FeCl2, concentrations above 5.5
mM supported maximal rates except at an 880 mM when the rate declined (Figure L3).
The ranges of molar ratios of 880 DMSD to H202 tested in this the rate declined ( Figure
[-3) The ranges of molar ratios of DMSD to H202 tested in this study were between 1:5
to 1:880. Only the higher molar ratios had an effect on mineralization. The maximal
initial rates were achieved at the ratios 1: 110 to 1:660. Pignatello (21) studied the effect
of H202 levels on 2,4-D mineralization by Fenton reagent and demonstrated that at ratios
of above 2021 H202:2,4-D, 2,4-D was mineralized to CO2 almost independently of H202
concentrations. The largest amount of H202 needed to mineralize DMSD as compared to

2,4-D is most likely the result of a higher degree of resistance of DMSD to oxidation.

35

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

100
so 1
so #9
70
so
50
40
a 30
I": 20
E 10 17°C?
a . .—
9?» t
.28 ‘40
#0
Foo,
I t .—<
20 25

Figure 1-3. Mineralization of DMSD in water by Fenton’s reagent with different
H202 concentrations (lmM DMSD, 2 mM iron catalyst). Treatments consisted of (I)
DMSD only, (0) iron only, (A ) 440 mM H202 only,( v) 880 mM H202, (O) 660 mM
H202, (+) 440 mM H202, ( X) 220 mM H202, (*) 110 mM H202, (O) 55 mM H202, (III)

5.5 mM H202

36

 

 

 

 

 

 

 

 

 

 

’6
(D
'0
‘0
rd
“5
O\0 .1 I I
‘6“ I l l =:
0 it i
a) .
as
0 3% air
:1
I .r I I I U H
10 15 20 25

Reaction time (hour)

Figure I—4. Mineralization of DMSD in water by Fenton’s reagent with different
DMSD concentrations (2 mM iron catalyst, 440 mM H202). Treatments consisted of (I)
0.1 mM DMSD only, (C) 3 mM DMSD only, (A) iron catalyst only,( v) H202 only, (6)

0.1 mM DMSD, (+) 1 mM DMSD, (X) 2 mM DMSD, (*) 3 mM DMSD.

37

DMSD mineralization in Soil Slum’es

A major advantage in working with Fenton reaction in soil is that soils naturally
contain sources of iron (e.g. iron oxyhydroxides) that can catalyze oxidation and
mineralization of the target contaminants. Unfortunately, soil colloidal material may

reduce the effectiveness of the Fenton oxidation by competing with target contaminants
for .OH radicals or catalyzing peroxide decomposition to H20 and O2 (19). Dissolved

natural organic matter can scavenge .OH radicals (59). Also, in contrast to water soluble
compounds, contaminants in the soils are usually sorbed to a certain extent or present as
nonaqueous phase liquids, and this can have significant effect on Fenton process
chemistry. Due. to its high water solubility DMSD is expected to stay mostly in solution
under ambient conditions in soil. Overall, the rates of DMSD mineralization were lower
for Fenton-reagent treated soil slurries than in water. This despite the fact that much
higher concentrations of iron (H) and H202 were used.

In the case of DMSD mineralization in soil only FeCl2 concentrations were
Optimized, as this is the classical Fenton reagent. The most effective level was found to
be 100 mM/kg (Figure I-5). Increasing the concentration above this level did not cause
increase in mineralization (data not shown). Control experiments ill soil with H202 and
no iron addition resulted in degradation of 10 % of the DMSD. Therefore, naturally
occurring iron in the form of amorphous and crystalline iron oxides (Table 1) may not

supply sufficient iron for Fenton-like reaction. These results are consistent with the

38

results of Watts et al. (20) who found that oxidation of sorbed pentachlorOphenol

was effective using low H202 concentration if catalysis was promoted by soluble iron

rather than natural iron hydroxides.

 

 

“C as (:02 (% of added)

 

 

 

 

 

 

 

 

 

Reaction time (day)

Figure I-5. Mineralization of DMSD in soil by Fenton’s reagent with different
iron catalysts concentrations (10 mM DMSD, 1.5 M H202). Treatments consisted of (I)

DMSD only, (+) H202 only, (A) 100 mM iron only,( v ) 50 mM iron only, (0) 100 mM
iron only, (X) 10 mM iron, (*) 25 mM iron, (A) 4 mM iron, (0) 50 mM iron, ([3) 75

mM iron, (0 )100 mM iron.

39

Seven iron catalysts were compared for their ability to catalyze DMSD
mineralization in soil (Figure I-6). Iron (HI) nitrate was most effective among those
tested. It promoted 42% oxidation of 10 mM DMSD in soil slurry. Other iron catalysts,
i.e. (iron (H) Chloride, iron (ID perchlorate, iron (H) sulfate, iron (H1)chloride, iron (H1)
perchlorate and iron (HI) sulfate) provided l6-32% DMSD mineralization under the same
reaction conditions. Iron (H) compounds were less effective catalysts than iron (HI)
species, presumably owing to the initial loss of H202 in oxidizing Fe (H) to Fe (H1). The
most effective iron catalysts were in the order:

N03' > so,” > C1‘ > cm;

The effectiveness Fe(NO3)3 is probably because of nitrate salts quenched .OH radicals-at
the lowest rates and, therefore, provided the highest degree of mineralization of DMSD.
Iron sulfates provided intermediate mineralization while perchlorates were least efficient
in promoting the mineralization. Control experiments (Figure I-6) with Fe (H)/Fe (HI)
alone showed no DMSD mineralization. Theinitial H202 concentration in soil affected
the initial rate and extent of DMSD mineralization (Figure I-7). Concentrations of H202
between 0.1 and 1.0 M were tested. The least effective level was 0.1 M. The most
effective level was 0.25 M. Increasing H202 level above 0.25 M did not increase the
initial rate or extent of DMSD mineralization. Concentrations of H202 over 2.0-2.5 M led
to violent reactions and hence were not evaluated further. This effective concentration is
moderate considering that several organic components of the system competed with
DMSD for oxidant, e.g. dissolved soil organic matter, and that soil particles can catalyze

H202 decomposition. Hydrogen peroxide is relatively unstable in soils and is affected by

40

the concentration of catalase, iron and manganese (21). In addition the pH of the soil-

peroxide slurry significantly affects H202 decomposition rates.

 

 

 

 

 

 

 

 

 

 

1“C as CO2 ("/o of added)

 

 

 

 

 

 

 

 

 

Reaction time (hour) _

Figure 1-6. Mineralization of DMSD in soil by Fenton’s reagent with different
iron catalysts (10 mg/kg DMSD, 100 mM Fe(H)/Fe(HI), 0.625 mM H202). Treatments
consisted of (a) Fenton’s reagent samples, (b) controls. (I) FeCl2, (O) FeCl3, (A)

Fe(304)2, (v ) 1:62.604)» (9) Fe(CIO4)2. (+) Fe(C104)3, ( X) Fe(NO3)3, (*) DMSD only.

(—) H202 only.

41

Watts et al. (46) found that H202 decomposed rapidly at pH>5, however at pH 3
the H202 consumption rates were low while the rate of contaminant degradation was
high. The higher H202 concentrations required to mineralize DMSD in soil versus water
were comparable to the results of previous Fenton-like soil remediation studies (14).
Watts et al. (20) found that 2 0.1 M H202 with soluble iron oxidized sorbed
pentachlorOphenol, and Spencer et al. (1) determined that 3 to 9 M H202 was required to
oxidize sorbed diesel using oxyhydroxides as the only source of iron catalysts. The use of
moderate concentrations of H202 (0.75 M) catalyzed by soluble iron appears to be the
most effective process conditions for the oxidation of DMSD. An attempt to increase the
mineralization of DMSD in soil Was done by adding a second increment H202 (0.65 M)
to the soil five hours after the initial Fenton reaction commenced (Figure I-8). The result
shows that the reaction can be driven further by additional H202 and a total increase of
about 10% was achieved. '

The extent of DMSD mineralization was also influenced by its initial
concentration in the soil. Figure I-9 shows that 1mg/kg was mineralized to the great
extent (about 35%). Increasing the concentration over 10 mg/kg resulted in a somewhat
lower extent of mineralization (29%).

The ability of Fenton-like reactions to remediate DMSD contaminated soil
increased as a function of the liquid volumes. The slurry volumes from 2.5 to 5 mL per
2.5 g soil were evaluated. The highest degree of mineralization was achieved in

solutionzsoil ratio 221 (Figure I-lO).

 

 

 

 

 

 

 

 

 

40
a a
c: at
.9 "‘
z: + F
s.—
O
o\°
ON
0
I!
I .Ti I i I I ##1
10 1 5 20 _ 25

Reaction time (hour)

Figure I-7. Mineralization of DMSD in soil by Fenton’s reagent with different
H202 concentrations (lOmg/kg DMSD, 100 mM iron catalyst). Treatments consisted of
(I) DMSD only, (0 ) iron only, (A) 0.1 M H202 only,( v ) 0.5 M H202 only, (0) 1.0 M

H202 only, (+) 0.1 M H202, (X) 0.25 M H202, (*) 0.5 M H202, (O) 0.75 M H202,

(D) l M H202.

43

 

451

 

 

14C as 002 (% od added)

 

 

 

 

 

o 5 10 15 * 20 25
Reaction time (hour)

Figure [-8. Mineralization of DMSD in soil by Fenton’s reagent with addition of
H202 at thr fifth hour (10 mg/kg DMSD, 100 mM iron (HI) catalyst, 0.625 M H202).

The vertical line indicate H202 addition. Treatments consisted of (I) DMSD only, (O)

Fer only, (A ) _Fe(ClOa)3 only, (v )Fe(N03)3 only, (+) FeCl3, (X) Fe(C104)3, (*)

FC(NO3)3 .

44

 

 

 

 

 

14C as CO2 (°/o of added)
"i"

 

 

 

 

 

 

 

Reaction time (hour)

Figure I-9. Mineralization of DMSD in soil by Fenton’s reagent with different
DMSD concentrations (100 mM iron catalyst, 0.625 M H202). Treatments consisted of
(I) 1 mg/kg DMSD only, (a) 100 mg/kg DMSD only, (A ) FeCl2 only,(v ) H202 only,
(6) 200 mg/kg DMSD, (+) 100 mg/kg DMSD, (X) 10 mg/kg DMSD, ( *) 1 mg/kg. I

DMSD.

45

 

 

 

 

 

 

 

 

4o 4
‘ X
A 35 - /X X
'0
Q)
.0
'8
“o r "‘
o\°
N
O
O
U)
(U .
:0 J., 5
[fl

 

 

10 15 20 25
Reaction time (hour)

Figure I-10. Mineralization of DMSD in soil by Fenton’s reagent at different
solutionzsoil rates (10 mg/kg DMSD, 100 mM FeC12, 2 M H202). Treatments consisted
of (I) 0.5/1 DMSD only, (C) 2/1 DMSD only, (A) 0.5/1 FeCl2 only,( v ) 2/1 FeCl2

only, (0) 0.5/1 H202 only, (+) 2/1 H202 only, (X) 2/1 Fenton’s reagent, ( *) l /1 Fenton’s

reagent, ([3) 0.5/l Fenton’s reagent.

46

SUMMARY

The characteristics of the Fe(H)/ H202 and Fe(IH)/H2O2 systems were evaluated
for the mineralization of DMSD in water and soil. It was found that the Fenton reagent '
could effectively mineralize DMSD in solution and soil slurries. The efficiency of the
Fenton reagent to destroy DMSD was affected by initial H202, iron and DMSD
concentrations as well as the oxidation state of the iron catalysts and iron counterions.
The rate of Fenton reaction was slowed in soil—water mixtures as Compared to pure water
and higher concentrations of the reagents had to be used. Mineralization of DMSD in
water increased when peroxide concentration increased from 5.5 to 110 mM, but
decreased slightly at a peroxide level to 880 mM. For the soil slurries a peroxide
concentration between 0.25 and 0.75 M was needed for maximal mineralization of
DMSD, generally between 32 and 42%. Also addition of a second increment of H202
improved mineralization by 10%. The time required for completion of aqueous phase
DMSD mineralization was 1 to 5 hours when FeCl3 was used while the Fe (H) system
needed 1 to 10 hours to reached the same extent of mineralization, ~ 72%. In the soil
slurry experiment the process of mineralization was completed somewhat faster,
generally<5 hours. This may be due to the faster decomposition of H202 catalyzed by the
soil particles. Seven iron species were investigated for their effectiveness in catalyzing
DMSD mineralization in aqueous and soil media. Generally iron (H) compounds were
less effective than iron (IH) due to the initial loss of H202 in oxidizing Fe(H) to Fe(HI);
There were differences in the initial rates of the mineralization process due to the

presence of different counterions in the solution. The most effective iron catalysts were

47

Fe(NO3)3 and FeCl3 in the aqueous phase and Fe(N03)3 in the soil slurries. A higher
percent of mineralization was achieved with lower initial DMSD concentration and
higher liquid to soil ratios.

The Fenton reagent effectively destroyed DMSD in aqueous solutions and
soil slurries, and indicated that Fenton oxidation may be a useful remediation treatment

for DMSD contaminated water and soils.

REFERENCES:

1. Buch, R.R.,' D.N.Ingerbrigsten, Rearragement of Poly(dimethylsiloxane) fluids on soil,

Environ.Sci.Technol., 6:676, 1979
2. Lehmann, R.G., S.Varaprath, C.L. Frye, Degradation of silicone polymers in soil,
Environ.Toxicol.Chem., 7: 1061, 1994
3. Carpenter, J.C., J.A.Cella, S.B.Dom, Study of degradation of PDMS on soil,
Environ.Sci. Technol., 292864, 1995
4. Barberini, M., C. Minero, E. Pelizzetti, E. Borgarello, N. Serpone, Chmosphere,
Chemical degradation of chlorophenols by F enton ’s reagent,l6:2225, 1987
5. Mohanty, N .R., I.W.Weil, Oxidation of 2,4-Dinitrotoluene using Fenton’s reagent.
Reaction mechnisms and their practical applications, Haz.Waste Haz. Mater., 102171,
1993
6. Seldak, D.L., A.W. Andren , Aqueous-phase oxidation of polychlorinated biphenyls

by hydroxyl radicals, Environ.Sci. Technol., 25: 1419, 1991

48

7. Watts, R.J., S.E.Dilly, Evaluation of iron catalysts for the Fenton-like remediation of
diesel contaminated soil, J.Haz.Mater., 51:209, 1996

8. Ruppert, G., R.Bauer, G.Heisler, The photo-Fenton reaction - an effective
photochemical wastewater treatment process, I.Photochem.Photobiol.A. Chem., 73:75;
1993

9. Gee, B.W., J.W.Bawder,. Particle-size analysis. in: Methods of soil analysis. Part 1
Physical and mineralogical methods, p.399-404, Madison, WI, 1986

10. Nelson, D.W., L.E. Sommers, Total Carbon, Organic Carbon, and Organic matter. in: '
Methods of soil analysis. Part 2 Chemical and microbiological Properties, p.539-579,
Madison, WI, 1986

11. Soil Conservation Service. Soil survey laboratory methods and procedures for
collecting soil samples, Soil survey investigatioaneport 1 US. Government Printing
Office. Washington DC, 1972

12. Jackson, M.L., C.H.Lin, L.W.Zelazny, in: Methods of soil analysis. Part 1 Physical
and mineralogical methods, p.113-124, Madison, WI, 1986

13. Ronen, Z., M.H. Gordon, J.M. Bollag, Biological and chemical minaralization of
pyridine, Environ.Sci.Technol., 13:21, 1994

14. Pignatello, J.J., Dark and photoassisted Fe(IH)—catalyzed degradation of
chlorophenoxy herbicides by hydrogen peroxide, Environ.Sci.Technol., 26:5, 1992

15. Kochany, J., E. Lipczinska—Kochany, Chemosphere, 25: 1769, 1992

16. Yang, M., J .Hu, K. Ito, Characterization of Fe(H)/H202/UV oxidation process,

Environ.Sci.Technol., 19:183, 1998

49

17. Silva, R.N., Rev.Pure Appl. Chem., 222115, IES Lighting handbook, 4th ed.,N.Y.,
1966

18. Anderson, L.I., J .0.johnson, R.F. Christman, Extent of ozone’s reaction with isolated
aquatic fulvic acid, Environ.Sci.Technol., 2:793, 1986

19. Schumb, W.C., C.N. Sherfield, R.L. Wentworth, Hydrogen peroxide, ACS
Monography, 128 Am.Chem.Soc., Washington DC, 1955

20. Watts, R.J.,M.D.Ude1, R.M.Monsen, Oxidation of pentachlorophenol catalyzed by
H202 and naturally occurring iron oxyhydroxides, Water Environ.Res., 65:893. 1993
21. Pignatello, J .J ., K.Baehr, Ferric complexes as catalysts for degradation of 2,4-D and

metolachlor in soil, J.Environ.Qual., 232355, 1994

50

CHAPTER II
INTRODUCTION

Polydimethylsiloxane (PDMS) is synthetically produced organosilicone
polymeric material widely used in industry and consumer products. PDMS is
characterized by low surface tension, resistance to chemical and thermal attack, water
repellency, and resistance to oxidation and UV reactions ( 1). Because of its recalcitrance,
many of the PDMS uses result in the presence of PDMS in waste waters. It can than enter
the environment as a component of effluents and sludges from wastewater treatment
plants. PDMS is expected to partition onto sediment solids and be applied to soil via land
application of sewage sludges (2).

Recent work by Lehman et al.(3) and Carpenter et al. (4) have shown that once
PDMS is introduced into Soil, it can undergo chemical hydrolysis to low molecular water-
soluble dimethylsilanediol (DMSD). The hydrolysis reaction is moisture dependent and
much faster in dry soil. Buch et al. (5) demonstrated the clay catalyzed conversion of
PDMS to water-soluble dimethylsiloxanoles and volatile oligomeric methylsiloxane
species. Recently Xu et a1. (6) found that soil factors such as clay content and clay type
are very important in determining the degradation rates of PDMS in soil. The differences
arrived from the different interaction strength between clay minerals and PDMS
degradation products. Surface area of the clay particles played an important role in

PDMS degradation. Higher surface area provided more PDMS degraded and speeded up

. the process.

51

 

Silicones may also be subject to photochemical degradation. Rapid degradation of
water soluble mixture of oligomeric dimethylsiloxanes by sun light in the presence of
environmentally relevant levels of nitrates has been reported (10). Silicate was
demonstrated to be the final photodegradation product of these water soluble siloxanes.
When an aqueous emulsion of high molecular weight PDMS was irradiated in the
presence of nitrate ion, a level of silicate corresponding to 1.7% of the original quantity
was found after 5 weeks. The low conversion was though to be due to the very low
solubility of PDMS such that only a monomolecular surface layer of the polymer could
react with the hydroxyl radicals formed in the water.

Fenton reaction (7) is recognized as one of the oldest powerful oxidizing reactions
available. This reaction has been used to decompose a wide range of refractory organic
pollutants (8, 9). The Fenton reagent is a mixture of H202 and ferrous salt which
produces hydroxyl radical ('OH). The ‘OH radical is capable of nonspecific oxidation of
many organic compounds.

The objective of the present research was to investigate the potential of the Fenton

reagent to mineralize PDMS in water and soil and to optimize the process conditions.

52

MATERIALS AND METHODS

PDMS-Water experiment

A stock solution of 0.97 mg/ml polydimethylsiloxane (PDMS) and ”C-PDMS:
dissolved in tetrahydrofurane (THF) was prepared in advance and placed into 60 ml

serum bottles. The total final PDMS concentration in water ranged between 0.05 to 200

mg/mL. The amount of '4 C-PDMS in each bottle was about 5 nCi. A stock solution of
20 mM iron (H) chloride was prepared fresh every time, and the concentrations evaluated
varied between of 2.0 to 3.0 mM. Hydrogen peroxide concentrations of 880 and 440 mM
were used to initiate the Fenton reaction. In several experiments the reaction reached
plateau due to the consumption of the hydroxyl radicals available. At this point H202
were added to the bottles, at the same initial concentrations, in order to stimulate the
mineralization process. Bottles were closed with rubber stoppers fitted with center well
(Kontes, Morton Grove, H.) containing a strip of 5-x-l cm filter paper and 0.2 ml 1N
NaOH used to trap l4CO2. Reactions were performed in dark with shaking at 100 rpm. At
specific intervals the filter paper was removed and placed in a vial with 10 ml
scintillation cocktail. Vials were left overnight in the dark and then radioactivity was
measured by liquid scintillation counting (LSC). Each treatment was performed in
duplicate, and all results are expressed as a mean value i standard error. Temperature
controlled experiments were performed in Control Environment Incubator Shaker (New
Brunswick Scientific, Edison, NJ). After the completion of the reaction 10 mL THF were
added to the reaction mixture to help dissolve the PDMS. The bottles were than shaken

overnight at 200 rpm. Aliquots (1 mL) were placed in 10 mL scintillation cocktail and

53

 

 

radioactivity remaining in the medium was measured; 80% of the radioactivity in the
aqueous phase was recovered by this method. Analysis of the degradation products are
subject of a future research.

The photo-Fenton reactions were carried out in Rayonet Photochemical Reactor (The
Southern New England Ultraviolet Co.) fitted with RUL 3000 lamps having a peak
energy output at 300 nm. During irradiation the samples were shaken at 200 rpm. The
irradiation chamber temperature was 30-40°C. The glass UV cut off was at 290 nm.

Cationic (cethyltrimethylammonium bromide), anionic (sodium dodecyl sulfate) and

 

neutral (Triton-X 450) surfactants were tested for their ability to increase PDMS
solubility in the water phase. They were dissolved in DI water and added to the bottles at

concentrations at, below and above their critical micelles concentration (CMC).

PDMS - soil experiments

PDMS dissolved in THF was applied to 2.5 g dry soil in 60 ml serum bottles in
the concentration range of 10- 1000 mg/kg. The maximum contaminant concentration was
based on typical PDMS levels found in wastewater treatment sludges (3). The THF was
allowed to evaporate completely. Aqueous solutions of iron (H) chloride and H202 were
added to soil. The concentrations of FeCl2 and H202 on a dry soil weight base were 100
mmol/kg and 1 to 2 mol/kg, respectively. Total liquid volume in the bottles was 2.5 mL.
The reactions were performed in the dark and bottles were shaken at 200 rpm.

All experiments were initiated by addition of H202 to the vials. Control
experiments using PDMS alone, combinations of PDMS and H202, and combinations of

PDMS and iron species were performed in parallel.

54

RESULTS AND DISSCUSION

Degradation of PDMS by Fenton reagent in water

PDMS is characterized by low surface tension, resistance to chemical and thermal
attack, water repellency, and resistance to UV radiation and oxidation. No data for its
biodegradation exist. The room temperature solubility of PDMS is very low. Therefore,
in mixtures with water ,essentially all PDMS existed as a separate liquid phase. PDMS
was presented as a 0.05 mg/L suspension in the reaction vessel. No evaporatlve loss of
PDMS from the vessel occurred.

Mineralization of PDMS under various reaction conditions was investigated ,
quantitating the 14CO2 released. The maximum observed conversion of PDMS to CO2
was 80% at 66°C in presence of 2mM FeCl2 and 880 mM H202 with stepwise addition of
H202 over 3 days (Figure H-l). Under similar conditions but at room temperature only
8.5% transformation was observed after 5 days (Figure H-2). Thus temperature plays an
important role in the Fenton degradation of PDMS. The effect of an incremental increase
in temperature from 42°C to 65°C to the extend of PDMS transformation is shown in
Figure H-3. At 42°C no '4 CO2 was trapped over a 2 day reaction time, indicating no
reaction. At 54°C the reaction commenced and was further accelerated at 66°C. A similar
temperature effect was observed in a study by Pignatallo and Chapa ( l l), where a two
phase mixture of comersial PCB oil Aroclor 1242 in water at 66°C was degraded by a
photoassisted Fenton-type reaction consisting of Fe(HI), H202, and UV light (300-400
nm).Up to 88% PCB removal and 85% dechlorination were obtained. Degradation at

room temperature was minimal. The reason for the improved degradation at 66 0C was

55

 

not clear according to the autors. They suggested that the temperature effect was due to
either enhanced rate of PCB dissolution into water, or inorganic phototchemical kinetic

factors affecting the oxidation reactions. An increase in the temperature leads to higher

rate of .OH formation and this may in part account for the higher reactivities observed by
PignatellO and Chapa, and in our studies with PDMS. In adition the improved degradation
at 66°C compared with room temperature may be due to the increased water solubility of
the PDMS at the higher temperature. It is reasonable to assume that little oxidation occurs
in the neat PDMS phase or for the PDMS absorbed or coated on surfaces.

Neither peroxide nor FeCl2 is expected to be appreciably soluble in the neat phase.
Increasing the solubility in water is likely to lead to partial degradation and formation of ,
lower molecular weight compounds. Such oxidized intermediates would be more soluble

than PDMS, thus increasing their effective reactivity in solution. Addition of H202 at this

point causes formation of reactive .OH radicals capable of fast degradation .of these

products.

56

 

 

 

 

 

 

 

 

 

 

 

 

80 - *
/X“"""/
X
A /
'o X
(I)
'o 50 -
(U ‘ *
E I /
\O x/ -—A
O
v 40 a
N
O
o ‘ I /
a) +
(U
20 -l
«O /
'- d / -—-/ —. . .
I / ———- -»——-—-~
I I ' I I I D I D T U I I I D
0 1 0 20 30 40 50 60 70 80

Reaction time (hour)

Figure H-l. Degradation of PDMS in water by Fenton’s reagent at 66°C and
addition of H202 indicated by the vertical lines (0.05 mg/L PDMS, 2 mM FeCl2).
Treatments consisted of (I) PDMS only, (0) 220 mM H202 only, (A) 880 mM H202

only,( v ) FeCl2 only, (0) 100 mM H202, (+) 220 mM H202, (X) 1 mM iron, (-) 440 mM

H202 , (*) 880 mM H202,

57

 

 

10

 

“c as co2 (% of added)

 

 

 

Reaction time (day)

Figure H—2. Degradation of PDMS in water by Fenton’s reagent and FeCl3/H2O2 at
room temperature (0.05 mg/L PDMS, 2 mM iron catalysts, 440 mM H202). Treatments
consisted of (I) PDMS only, (0) H202 only, (A) FeCl2 only,( v ) FeCl3 only, (0)

Fenton’s reagent, (*) FeC13/H202,

58

 

30
‘ 42°C 46°C 54°C 66°C 1
25 d

1“c as co2 (% of added)

 

 

 

 

 

 

 

Reaction time (day) I

Figure H—3. Degradation of PDMS in water by Fenton’s reagent and at different
temperatures and adddition of H202 indicated by the arrow. Treatments consisted of (I)

0.05 mg/L PDMS, 2 mM FeCl2, 440 mM H202 , (O ) 0.05 mg/L PDMS, 2 mM FeC12,

880 mM H202, (A) 0.05 mg/L PDMS, 3 mM FeCl2, 440 mM H2O2,( v ) 0.05 mg/L

PDMS, 3 mM FeCl2, 880 mM H202, (O) FeCl2 only.

59

 

 

 

35

 

 

 

 

 

 

" 3&1-
.0
(D
:8 2.5% 3*
as o/,
*5 -/
O 2.0‘ /
o\
301 d /*

L5“
0
8 .

LO- ’flflflfl,,,.»~vf::::::z
o ”#2,..-
v I

0.54 / M/‘T"

. / //.._. s
o .//""
0.0 l 1 I 1 l I l I
0 5 10 15 20 25

Reaction time (hour)

Figure H-4. Degradation of PDMS in water by Fenton’s reagent at room
temperature and different PDMS concentrations ( 2 mM FeCl2 , 440 mM H202 ).
Treatments consisted of (I) PDMS only, (0) H202 only, (A) FeCl2 only,( v ) 0.01

mg/ml PDMS, (0) 0.1 mg/ml PDMS, (*) 0.2 mg/ml PDMS.

60

The extent of PDMS degradation depends on its initial concentrations in the

aqueous phase. The higher concentrations were oxidized faster and to a higher extent at
room temperature (Figure H—4). However, at 66°C 0,05 mg/L PDMS was oxidized to a

greater extent than 0.1 mg/L PDMS (Figure H-S).

 

70

 

60 -‘ -_.———/

50-

 

 

1“c as co2 (% of added)

 

 

 

 

 

 

o 10 20 so 40 so
Reaction time (hour)

Figure H-5. Degradation of PDMS in water by Fenton’s reagent at different
PDMS concentrations and addition of H202 indicated by the vertical line ( 2 mM FeCl2 ,
880 mM H202 ,66°C). Treatments consisted of (I) 0.1 mg/L PDMS and H202 only, ( O)

’ 0.05 mg/L PDMS and H202 only, (- ) 0.1 mg/L PDMS (v ) 0.05 mg/L PDMS.

61

 

Tetrahydrofurane (T HF) had a detrimental effect on the Fenton reaction. Addition
of 0.1 ml PDMS dissolved in THF (without evaporation of the THF) lead to 20 % less

transformation of the PDMS as compared to the case where THF was evaporated (Figure

H-6). Apparently, THF is quenching .OH radical formation and should be completely

evaporated prior to addition of the other Fenton system components.

 

 

 

 

25
j /’/.
20.. . /
6 c
Q) -I
'C
U
(U 15-
u—
0
52 J
N 1o-
0 o
O ‘ /
U)
(u ///.
_____._.._./f'
0L I ‘ I ' I I I

 

o 10 20 so 40 so
Reaction time (hour)

Figure H-6. Degradation of PDMS in water by Fenton’s reagent at 66°C
and adddition of H202 indicated by the vertical line with and without THF ( 0.05 mg/L
PDMS, 2 mM FeCl2, 440 mM H202). Treatments consisted of (I) PDMS, (O) PDMS +

THF.

62

The effect of Fe(II) compared with Fe (III) at 66°C is negligible. Both iron

catalysts resulted in about 70% mineralization after 70h (Figure H-7) and can be used

 

 

 

 

 

 

 

interchangeable.
80
70 -
i
{a 60 -
(D .
38
CD 50 -
B a
°\o 4O -
N "i
O 30 -
0 .
3 20 -
O . -
:1
10 -
o f i F I ' 1 l 1 1 I I ‘I l V r
0 1 0 20 30 4O 50 60 70

Reaction time ( hour)

Figure II-7. Degradation of PDMS in water by Fenton’s reagent and by Fer/
H202 at 66°C and adddition of H202 indicated by the vertical line ( 0.05 mg/L PDMS, 2

mM iron catalyst, 440 mM H202). Treatments consisted of (I) Fe (H), (O ) Fe(IH).

63

Three types of surfactants were added in an attempt to increase PDMS solubility
and hence increase the rate of reaction at room temperature. Cationic
(cethyltrimethylammonium bromide), anionic (sodium dodecyl sulfate) and neutral
(Triton-X 450) surfactants at concentrations below and above their critical micelle
concentration (CMC), as well as very low concentrations (0.01 mg/ml) were evaluated.

All the surfactants tested had a strong detrimental effect on the PDMS degradation. They

probably served as competing substrates and undergo oxidation by .OH radicals instead
of PDMS, and/or they may have complexed Fe(H) leading to its inactivation.

An attempt to improve the degradation of PDMS was done by using a photo-
Fenton oxidation system at room temperature. Irradiation of an aqueous solutions (0.05
and 0.1 mg/L) of PDMS with UV resulted in no greater formation of '4 C02 as
compared to non-irradiated controls. This result is consistent with a study (10) that has
shown only 1.7 % degradation of PDMS after 37 days of irradiation with UV light in
presence of nitrate ions. Hydroxyl radicals are formed during irradiation of nitrate or
nitrite ions (10). The slow degradation was thought to be due to the very low water

solubility of PDMS and the limited surface area of the polymer available to react with the

‘OH radicals.

Degradation of PDMS by Fenton reagent in soil

PDMS is expected to partition onto solids and be sorbed by the soil due to its low
water solubility. This fact, combined with the potential of competition of soil colloidal
materials for Fenton reactive intermaediates, was expected a priori to reduce the

conversion of PDMS to C02. When the reaction was conducted at room temperature for

64

24 hours a maximum conversion of only about 3% PDMS to C02 was obtained. This is
less than half the maximum conversion obtained in aqueous (soil-free) systems at room
temperature. The process was influenced by the initial PDMS concentrations (Figure 11-
8); lower concentrations were degraded to a higher extent.

As in pure water, temperature also was a major factor that influenced the
Fenton oxidation of PDMS in soil. Figure II—9 shows the effect of a step wise increase in
the temperature from 46°C to 68°C at different H202 concentrations. The highest
degradation of 12% was achieved at 68 °C and 2 mol/kg H202. In attempt to increase the
minaralization of PDMS, H202 was added incrementaly. The highest percent degradation
achieved was 20% for 7 days at 66°C, during which, portions of H202 were added 3 times
(Figure 11-10). The control which contained PDMS and FeClz had 14 % 1“ co2 trapped
indicating that volatilization is probably a major route of the polymer transformation.
This phenomena may be explained with the catalytic role of the soil clay minarals, which

can promote PDMS degradation.

65

 

1&5

 

 

 

 

 

 

 

 

 

 

 

1104
____’_______._———O
______'_____...
A “’10”.
.0 //
8 2.5 o[/0
“O
(U
75 2L0-
\o :1
O
‘1. A
8 /://v
If...
(D
(U
0
3 ,5;
/ a
I fl I f
15 20 25

 

Reaction time (hour)

Figure II-8. Degradation of PDMS in soil by Fenton’s reagent at room temperature
and different PDMS concentrations ( 100 mg/kg FeClz , 0.7 M H202 ). Treatments
consisted of (I) 10 mg/kg PDMS only, (0) 100 mg/kg PDMS only, (A ) 500 mg/kg
PDMS only, (v) 1000 mg/kg PDMS only, (A) H202 only, ( *) FeClz only, (0) 10 mg/kg
PDMS Fenton, (Cl) 100 mg/kg PDMS Fenton, (0)500 mg/ltg PDMS Fenton, (V)lQOO

mg/kg PDMS Fenton.

66

 

 

12
. 46°C 54°C 65°C 68°C v
/
10 - v///
/

 

 

“c as 002( % of added)

 

 

 

 

 

 

: I ' I ' I ' I I I I
0 2 4 6 8 10 12 14 16

Reaction time ( day)

Figure lI-9. Degradation of PDMS in soil by Fenton’s reagent at diferent
temperatures and different H202 concentrations ( 10 mg/kg PDMS, 100 mM FeClz ).

Treatments consisted of (I) 0.5 mM H202, (O) 0.75 mM H202, (A ) 0.5 mM H202,

(')2mMH202.

67

 

 

 

 

 

20
6 15-
(D
8 v a
(U
s..—
0
°\0 10-
N
8 /I
8 5-
.0
" J
0 I I T I I Tl T f T

 

 

 

 

' I I I '
‘0 20 4O 60 80 100 120 140 160 180

Reaction time (hour)

Figure II-lO. Degradation of PDMS in soil by Fenton’s reagent at 66°C and
addition of H202 at different concentrations ( 10 mg/kg PDMS, 100 mM FeClz ).

Treatments consisted of (I) 1.0 mM H202 only, (0) 2.0 mM H202 only, (A) 1.0 mM

H202, (v ) 2 mM H202 , (.) PDMS + FBClz only.

68

 

CONCLUSIONS
In summary, PDMS are oxidized in aqueous systems containing ferrous ion and
hydrogen peroxide above room temperature. PDMS degradation of 80% in aqueous
system and 20 % in soil was achieved by repeated addition of H202 and 66°C.
Temperature is obviously a critical variable. After further development these systems

might be useful for treatment of wastewaters and soils contaminated with PDMS.

REFERENCES
1. Noll, W., Chemistry and Technology of Silicones, Academic Press Inc., New York, NY
10003, 1968
2. Lentz, C., Silicones in the environment, Ind.Res.Dev., 222139, 1980
3. Lehmann, R.G., S.Varaprath, C.L. Frye, Degradation of silicone polymers in soil,
Environ.Toxicol.Chem., 7: 1061 , 1994
4. Carpenter, J.C., J.A.Cella, S.B.Dom, Study of degradation of PDMS on soil,
Environ.Sci. Technol., 292864, 1995
5. Buch, R.R., D.N.Ingerbrigsten, Rearragement of Poly(dimethylsiloxane)fluids on soil,
Environ.Sci.Technol., 6:676, 1979
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7. Fenton, H.J.H., On a new reaction of tartaric acid, Chem.News, 28:39, 1876
8. Mohanty, N.R., I.W.Weil, Oxidation of 2,4-Dinitrotoluene using Fenton ’s reagent.
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1993

69

9. Sun, Y., J .1 . Pignatello, Photochemical reactions involved in the total mineralization of
2,44) by Fe3*/Hzo,/UV, Environ.Sci.Technol., 27:304, 1993
10. Buch, R.R., T.H.Lane, R.B.Annelin, C.L.Frye, Photolytic oxidative demethylation of

aqueous dimethylsiloxanols, Environ.Toxicol.Chem., 6:676, 1984

ll. Pignatello, J .J ., G.Chapa, Degradation of PCBs by ferric ion, H202 and UV light,
Environ.Toxicol.Chem., 13:423, 1994

70