Degradation of Hydrophilic and Hydrophobic

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DEGRADATION OF HYDROPHILIC AND HYDROPHOBIC

CONTAMINANTS IN PERCARBONATE SYSTEMS

SATTAM FAHAD ALMOJIL

A dissertation submitted in partial fulfillment of the requirements for the degree of

DOCTOR OF PHILOSOPHY

WASHINGTON STATE UNIVERSITY Department of Civil and Environmental Engineering

DECEMBER 2015

To the Faculty of Washington State University:

The members of the Committee appointed to examine the dissertation of SATTAM

FAHAD ALMOJIL find it satisfactory and recommend that it be accepted.

______Richard Watts, Ph.D., Chair

______Amy Teel, Ph.D.

______I. Francis Cheng, Ph.D.

______Indranil Chowdhury, Ph.D.

ACKNOWLEDGMENT

I am really thankful and grateful for my advisor and committee chair, Dr. Richard Watts for his ideas, guidance, and support. I would like to thank my committee members: Dr. Amy Teel, Dr.

Francis Cheng, and Dr. Indranil Chowdhurt for their time and input. I am really thankful for Dr.

Amy Teel for her help to improve my writing and also to use KaleidaGraph program. I would like to thank my previous lab mates Dr. Mushtaque Ahmad, Dr. Marissa Merker, Dr. Farah Elloy,

Dr. Miao Yu, and Syeed Iskandar for answering my questions.

Acknowledgement would not be complete without mentioning my parents, wife, brothers, sisters, and friends for motivating and supporting me. My special appreciation for my wife for her understanding and love during my graduate study at WSU. At the end, I would like to thank the ministry of education for providing me a scholarship to study M.S. and Ph.D. programs in the

United States of America.

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DEGRADATION OF HYDROPHILIC AND HYDROPHOBIC

CONTAMINANTS IN PERCARBONATE SYSTEMS

Abstract

by Sattam Fahad Almojil, Ph.D. Washington State University December 2015

Chair: Richard Watts

Three topics related to the degradation of hydrophobic and hydrophilic compounds in percarbonate systems were studied. The first chapter provides a literature review and a brief description of the three treatment systems used in the present research: catalyzed hydrogen peroxide propagations (CHP), sodium percarbonate (PC), and base-activated persulfate (BAP) treatment systems. The second chapter presents a study of the degradation of contaminants by superoxide generated in PC systems, and further investigates the relationship between type of halogen and degree of halogenation in superoxide degradation rate for halogenated contaminants.

Superoxide was responsible for the degradation of chlorinated compounds; and the rate of degradation by superoxide increased with greater degree of halogenation. Moreover, compounds with higher electronegativity were degraded more rapidly than those with lower electronegativity.

Degradation of hydrophobic contaminants sorbed onto silica gel in PC, CHP, and BAP treatments are discussed in chapter three. The results demonstrate that chloroalkene and chloroalkane compounds were effectively degraded in CHP and PC systems, but not in BAP

systems. Organochlorine insecticides and polychlorinated biphenyls were degraded in the CHP and BAP systems, but not in the PC systems. The hydrophobic herbicides diuron, bentazon, 2,4-

D, and 2,4,5-T were degraded in CHP, BAP, and PC treatment systems with a degradation rate above the desorption rate of gas purge (GP) systems. However, atrazine was degraded only in

CHP and BAP systems, indicating that superoxide does not react with atrazine.

The fourth chapter covers changes in superoxide reactivity in aqueous systems with the addition of solids and minerals. The effects of 14 minerals and solids on the reactivity of superoxide in

PC systems were investigated. The results showed increased superoxide reactivity in aqueous superoxide systems with the addition of solids or minerals due to increased surface area in the system. The results indicate that increased surface area—not the nature of the solid or mineral— is responsible for the increase in superoxide reactivity in aqueous PC systems. The linear relationship between the rate constants for HCA degradation and the actual solid or mineral surface areas demonstrates that providing surfaces in superoxide systems enhances its reactivity.

TABLE OF CONTENTS Page

ACKNOWLEDGEMENT ...... iii

ABSTRACT ...... iv

TABLE OF CONTENTS ...... vi

LIST OF TABLES ...... viii

LIST OF FIGURES ...... ix

CHAPTER 1: INTRODUCTION Catalyzed Hydrogen Peroxide Propagation (CHP) Chemistry ...... 3 Persulfate Chemistry ...... 7 Sodium Percarbonate Chemistry...... 12 Objectives ...... 13 References ...... 16

CHAPTER 2: DEGRADATION OF CONTAMINANTS WITH SUPEROXIDE AND THE RELATIONSHIP BETWEEN DEGREE OF HALOGENTATION AND THE TYPE OF HALOGEN IN THE RATE OF ATTACK BY SUPEROXIDE GENERATED IN PERCARBONATE SYSTEM Introduction ...... 22 Methodology ...... 25 Results and Discussion ...... 28 Conclusion ...... 32 References ...... 34 List of Figures ...... 38

CHAPTER 3: DEGRADATION OF HYDROPHOBIC CONTAMINANTS SORBED ONTO SILICA GEL AND COMPARING PERCARBONATE TREATMENT TO CHP TREATMENT AND BASE-ACTIVATED PERSULFATE TREATMENT Introduction ...... 47 Methodology ...... 50 Results and Discussion ...... 53 Conclusion ...... 61 References ...... 63

List of Figures ...... 67

CHAPTER 4: INCREASING REACTIVITY OF SUPEROXIDE WITH SOLIDS AND SALTS Introduction ...... 87 Methodology ...... 88 Results and Discussion ...... 90 Conclusion ...... 94 References ...... 95 List of Figures ...... 98

vii

LIST OF TABLES Page

CHAPTER 2: DEGRADATION OF CONTAMINANTS WITH SUPEROXIDE AND THE RELATIONSHIP BETWEEN DEGREE OF HALOGENTATION AND THE TYPE OF HALOGEN IN THE RATE OF ATTACK BY SUPEROXIDE GENERATED IN PERCARBONATE SYSTEM Table 1. Groups of contaminants and their derivatives ...... 37

CHAPTER 3: DEGRADATION OF HYDROPHOBIC CONTAMINANTS SORBED ONTO SILICA GEL AND COMPARING PERCARBONATE TREATMENT TO CHP TREATMENT AND BASE-ACTIVATED PERSULFATE TREATMENT Table 1. Selected groups of hydrophobic contaminants ...... 66

CHAPTER 4: INCREASING REACTIVITY OF SUPEROXIDE WITH SOLIDS AND SALTS Table 1. Minerals formulas and surface areas ...... 97 Table 2. Mass (g) and surface area (m2) of added solids and minerals ...... 97

viii

LIST OF FIGURES Page

CHAPTER 2: DEGRADATION OF CONTAMINANTS WITH SUPEROXIDE AND THE RELATIONSHIP BETWEEN DEGREE OF HALOGENTATION AND THE TYPE OF HALOGEN IN THE RATE OF ATTACK BY SUPEROXIDE GENERATED IN PERCARBONATE SYSTEM Figure 1; Degradation of 25 µM dichloromethane, CF and CT in 0.5 M PC systems ...... 39 Figure 2; Degradation of 15 µM 1,2-DCA, 1,1,1-TCA, and PCA in 0.5 M PC systems .40

Figure 3; Degradation of 10 µM 1,2-DCE, TCE, and PCE in 0.5 M PC systems ...... 41

Figure 4; Degradation of 2 µM of chlorobenzene, 1,2-dichlorobenzene, 1,2,4 trichlorobenzene,1,2,3,4-tetrachlorobenzene, and pentachlorobenzene in 0.5 M PC systems ...... 42 Figure 5; Degradation of 10 µM trichlorofluoromethane, CT, and bromotrichloromethane in 0.5 M PC systems...... 43

Figure 6; The average percentage of H2O2 during the degradation processes for methane, ethane, ethylene, and chlorobenzene derivatives in 0.5 M PC systems 44 Figure 7; Degradation of 2 mM nitrobenzene, 1,3-dinitrobenzene, and 1,3,5 trinitrobenzene in 0.5 M PC systems ...... 45 Figure 8; Degradation of 5 mM 1,4-dioxane, MEK, 1-pentanol, MTBE and toluene in 0.5 M PC systems ...... 46

CHAPTER 3: DEGRADATION OF HYDROPHOBIC CONTAMINANTS SORBED ONTO SILICA GEL AND COMPARING PERCARBONATE TREATMENT TO CHP TREATMENT AND BASE-ACTIVATED PERSULFATE TREATMENT Figure 1; Degradation of 1,2-dibromotetrachloroethane in CHP, BAP, and PC systems ..68 Figure 2; Degradation of HCA in CHP, BAP, and PC systems ...... 69

Figure 3; Degradation of 1,3-dichloropropene in CHP, BAP, and PC systems ...... 70

Figure 4; Degradation of hexachloropropene in CHP, BAP, and PC systems ...... 71

Figure 5; Degradation of hexachloro-1 3-butadiene in CHP, BAP, and PC systems ...... 72

Figure 6; Degradation of hexachlorocyclopentadiene in CHP, BAP, and PC systems ....73

Figure 7; Degradation of DDE in CHP, BAP, and PC systems ...... 74

Figure 8; Degradation of DDD in CHP, BAP, and PC systems ...... 75

Figure 9; Degradation of lindane in CHP, BAP, and PC systems ...... 76

Figure 10; Degradation of DDT in CHP, BAP, and PC systems ...... 77 ix

Figure 11; Degradation of 2-chlorobiphenyl in CHP, BAP, and PC systems ...... 78

Figure 12; Degradation of 2,3-dichlorobiphenyl in CHP, BAP, and PC systems ...... 79 Figure 13; Degradation of 2,2',3-trichlorobiphenyl in CHP, BAP, and PC system ...... 80

Figure 14; Degradation of 2,2',3,3'-tetrachlorobiphenyl in CHP, BAP, and PC systems .81

Figure 15; Degradation of diuron in CHP, BAP, and PC systems ...... 82

Figure 16; Degradation of atrazine in CHP, BAP, and PC systems ...... 83

Figure 17; Degradation of bentazon in CHP, BAP, and PC system ...... 84

Figure 18; Degradation of 2,4-dichlorophenoxyacetic acid in CHP, BAP, and PC systems ...... 85

Figure 19; Degradation of 2,4,5-trichlorophenoxyacetic acid in CHP, BAP, and PC systems ...... 86

CHAPTER 4: INCREASING REACTIVITY OF SUPEROXIDE DUE TO THE PRESENCE OF SOLIDS AND SALTS Figure 1; Effect of adding 3 g glass beads (GB) on degradation of 2 μM HCA in 0.25 M PC systems ...... 99

Figure 2; Effect of adding 0.1 g iron minerals on degradation of 2 μM HCA in 0.25 M PC systems ...... 100

Figure 3; Effect of adding 0.1 g pyrolusite and manganite and 0.03 g birnessite on degradation of 2 μM HCA in 0.25 M PC systems ...... 101

Figure 4; Effects of adding 0.1 g trace minerals on degradation of 2 μM HCA in 0.25 M PC systems ...... 102

Figure 5; Effects of adding 0.05 g trace minerals on degradation of 2 μM HCA in 0.25 M PC systems ...... 103

Figure 6; Relationships between the actual surface areas of glass beads (GB) and minerals and rate constants for degradation of 2 μM HCA in 0.25 M PC systems ...... 104

CHAPTER 1 Introduction

Increasing levels of industrial and agricultural activities around the world have generated considerable amounts of hazardous organic compounds. Soil and water contamination from these compounds damage the environment and can cause diseases such as cancer. Organochlorine pesticides, for example, are widely used as insecticides and are known to accumulate in soils, plants, and animals.

A number of technologies have been investigated for soil and water remediation over the past thirty years. As with all technologies, they have limitations. Methods of treating contaminated soil and groundwater are divided into two types: ex situ and in situ. In ex situ remediation, contaminated media is removed and disposed of or treated before being returned to the source.

Soil is generally removed through excavation, while groundwater is removed by pumping. Watts and Teel (2005) note that pump-and-treat systems have been widely used over the past thirty years, but are now considered ineffective due to slow rates of contaminant dissolution and desorption.

In situ remediation is accomplished without removing contaminated soil or groundwater and is therefore less expensive than ex situ remediation. Arctander and Bardos (2002) suggested that while in situ remediation is more cost-effective than ex situ, effectiveness is dependent upon contaminant volume and distribution. The most common in situ remediation technologies include bioremediation, soil vapor extraction, and chemical oxidation. In bioremediation, oxygen is injected into the subsurface to enhance biological degradation activity.

Bioremediation is widely used to degrade many types of contaminants. A notable exception is sorbed contaminants (hydrophobic compounds) which must desorb before they can be treated through biological metabolism (Aronstein et al., 1991). Chemical oxidation is applied by injecting reactive chemical oxidants, such as potassium permanganate, ozone, or hydrogen peroxide, into the subsurface (Watts & Teel, 2006). In situ chemical oxidation (ISCO) is an increasingly common technology for soil and groundwater remediation. It is faster and more efficient than natural attenuation or bioremediation. According to Smith et al. (2006), one third of the current remediation technologies rely on ISCO processes.

The most popular oxidants used for ISCO include permanganate (MnO4ˉ), catalyzed hydrogen

2 peroxide propagations (CHP), ozone (O3), and persulfate (S2O8 ˉ) (Watts & Teel, 2006). Each of these oxidants has diverse reactivities with contaminants. Permanganate oxidizes contaminants through electron transfer. Although permanganate reacts rapidly with alkenes, it is incapable of degrading recalcitrant compounds such as carbon tetrachloride (CT), chloroform (CF), methylene chloride (MC), benzene, chlorobenzene, some pesticides, 1,1-dichloroethane (DCA),

1,1,1-trichloroethane (TCA), and PCBs (Huling and Pivetz, 2006). Catalyzed hydrogen peroxide reacts with a wide range of contaminants of concern, and it is also applicable for desorption of sorbed contaminants (Watts and Teel, 2005; Corbin et al., 2007). Furthermore, due to the generation of superoxide, CHP can dissolve and destroy dense non-aqueous phase liquids

(DNAPLs) (Teel et al., 2002; Smith et al., 2006). Ozone is also a strong oxidant that reacts with numerous contaminants. Persulfate is the newest ISCO oxidant. It is considered more reactive than permanganate and more stable than CHP (Liang et al., 2003).

Each of the ISCO reagents has one or more limitations in stability or transport. Ozone is limited by its low solubility in groundwater and by bubble distribution (Huling and Pivetz, 2006).

Permanganate is stable in the subsurface and persists for months (Siegrist et al., 2001); however, in soils with a high level of natural organic matter (NOM) permanganate is rapidly consumed, decreasing its transport to contaminated areas (Watts, 2011). Catalyzed hydrogen peroxide transport is limited by its stability in the subsurface. Chen et al. (2001) demonstrated that hydrogen peroxide, the oxidant reagent for CHP, is unstable in the subsurface. Khan and Watts

(1996) found that the half-life of hydrogen peroxide is no more than 48 hours, which negatively affects the transport of the oxidant species to contaminated regions. Persulfate is highly stable, persists for weeks, and is transported through porous media (Huang et al., 2002; Liang et al.,

2003). It is reactive with most contaminants when activated, but does not provide enhanced treatment of sorbed contaminants and the cost of persulfate is greater than permanganate and hydrogen peroxide (Teel et al, 2008). A new possible ISCO reagent is sodium percarbonate (PC).

It is also called solid hydrogen peroxide (H2O2) since it contains about 30% H2O2. Several studies have shown that PC has the potential for degradation of organic compounds.

Catalyzed H2O2 Propagations (CHP) Chemistry

CHP is a modification of the standard Fenton’s reaction. The standard Fenton’s reagent is prepared by adding dilute hydrogen peroxide (H2O2) to a solution of iron (II) at a pH of approximately 3, producing hydroxyl radical (OH•) (Haber & Weiss, 1934):

+2 • +3 H2O2 + Fe → OH + Fe (1)

Hydroxyl radical is a strong oxidant that reacts quickly with more than 98% of organic compounds at near diffusion-controlled rates in the aqueous phase (Haag & Yao, 1992).

However, its reactivity becomes negligible in the non-aqueous phase (Sheldon & Kochi, 1981) and is not available to react with sorbed contaminants or non-aqueous phase liquids (NAPLs)

(Sedlak & Andren, 1994; Watts et al., 1999).

Various methods are used to apply CHP, including increasing hydrogen peroxide concentration up to 1–5 M (Watts et al., 1990). Another method is to add different catalysts such as iron pyrophosphate (Wang & Brusseau, 1998), iron chelates (Pignatello & Baehr, 1994; Sun &

Pignatello, 1992), and iron oxyhydroxides (Miller & Valentine, 1994; Ravikumar & Gurol, 1994;

Tyre et al., 1991; Watts et al., 1993, 1997). As a result of increasing concentrations of hydrogen peroxide, a series of reactions are promoted to provide reactive oxygen species such as

• • perhydroxyl radical (HO2 ), hydroperoxide anion (HO2ˉ), and superoxide radical anion (O2 ˉ)

(Watts et al., 1990; Barb et al., 1951; Walling, 1975; De Laat & Gallard, 1999). These reactions are as follows:

• • H2O2 + OH → HO2 + H2O2 (2)

+3 • + +2 H2O2 + Fe → HO2 + H + Fe (3)

• • + HO2 ↔ O2 ˉ + H pKa= 4.8 (4)

• +2 +3 HO2 + Fe → Fe + HO2ˉ (5)

• Hydroperoxide anion (HO2ˉ) is a strong nucleophile. Perhydroxyl radical (HO2 ) is a weak oxidant that is unreactive with most organic compounds; however, it contributes to generating superoxide in CHP reactions (Reaction 4). When perhydroxyl radical is oxidized, hydroperoxide anion is formed (Reaction 5). McKillop and Sanderson (2000) showed that the electrophilic form of an oxidant is usually more active for oxidation in acidic conditions. In contrast, nucleophilic

forms of oxygen are more active under basic conditions. Superoxide is a weak nucleophile and reductant that has been considered a reactive species for the degradation of highly oxidized organics such as carbon tetrachloride (Smith et al., 2004; Teel & Watts, 2002). Nucleophile species such as superoxide and hydroperoxide anion degrade some organic compounds by reacting with electron poor moieties such as carbonyl carbon (David & Seiber, 1999); they also degrade halogenated compounds such as carbon tetrachloride (CT) and trichloroethylene (TCE).

In CHP systems, superoxide is identified as the species responsible for reduction of sorbed metals that are sorbed by kaolinite (Monahan et al., 2005). The reactivity and lifetime of superoxide in deionized water are much shorter than in organic solvents because superoxide is strongly solvated by water molecules (Afanas’ev, 1989). The reactivity of superoxide in the aqueous phase can be increased by adding solvents, such as H2O2, which are less polar than water (Smith et al., 2004). Several studies have demonstrated that superoxide is an applicable species for enhanced degradation of organic chlorinated contaminants, desorption of sorbed contaminants, and destruction of DNAPL organic contaminants.

Teel and Watts (2002) demonstrated that carbon tetrachloride (CT) is not reactive with hydroxyl radical in the traditional Fenton’s reaction, and they proved that degradation of CT and some other highly oxidized chlorinated organic compounds (e.g., hexachloroethane, bromotrichloromethane) in CHP reactions is mediated by superoxide (Teel & Watts, 2002).

Adding solvents to aqueous systems increases the reactivity of superoxide to degrade highly chlorinated compounds (Furman et al., 2009). Therefore, in the aqueous phase, adding solvents increases the reactivity of superoxide; the reactivity increases with solvent concentration.

According to Smith et al. (2004), degradation of carbon tetrachloride in CHP systems is

enhanced by increasing the concentration of H2O2 in water (Smith et al., 2004). In addition, Xu et al. (2010) pointed out that superoxide anion, which is generated in bicarbonate-activated H2O2

(BAP) systems by bicarbonate anion, is considered a nucleophilic species that degrades organic contaminants in wastewater. Finally, the results demonstrate the potential for superoxide to enhance the degradation and transformation of chlorinated compounds in soil and water.

Organic hydrophobic compounds that are sorbed are difficult to treat. The transformation rate of sorbed organic compounds at a contaminated site is usually related directly to the rate of desorption. Therefore, transformation occurs after sorbed contaminants are desorbed into the aqueous phase. However, recent results demonstrate that enhanced desorption of sorbed contaminants is promoted by non-hydroxyl radical species. Although many organic compounds are oxidized effectively by hydroxyl radical, Sedlak and Andren (1991) found that polychlorinated biphenyls (PCBs) that were sorbed to diatomaceous earth were not oxidized by hydroxyl radical. On the other hand, using high concentrations of hydrogen peroxide, the standard Fenton’s reagent can be modified to CHP, which enhances the treatment of sorbed compounds that are not desorbed by hydroxyl radical. Watts et al. (1994) found that in CHP reactions (Reactions 2-4), when H2O2 ≥ 3.4% (100 mM), degradation of hexachlorobenzene was faster than its loss by gas-purge desorption, which is a measurement of the maximum natural rate of desorption. Furthermore, Watts et al. (1997) reported that a group of chlorobenzene compounds that were sorbed by hematite were desorbed rapidly by CHP reactions. Additionally,

Corbin et al. (2007) found that superoxide anion, which is generated in CHP, is the species responsible for enhanced desorption of dodecane sorbed onto silica sand.

Dense non-aqueous phase liquids are also a significant remediation challenge. Because these contaminants are located deep in the subsurface, they are hard to detect and treat. For example, most reactants, such as hydroxyl radical, do not react in the non-aqueous phase. Sheldon and

Kochi (1981) reported that hydroxyl radical is only reactive in the aqueous phase. Previous studies show many difficulties with the destruction and transformation of DNAPLs. The most common contaminants in DNAPLs are TCE, perchloroethylene (PCE), and CT. According to

Stroo et al. (2003), many factors affect the destruction of DNAPLs, such as the rate of dissolution of contaminants, the surface to volume ratio of DNAPL, and the rates of reaction of the contaminants with the reactive oxygen species. In an investigation of which reactive oxygen species in CHP reactions are responsible for the destruction of carbon tetrachloride DNAPLs,

Smith et al. (2006) reported that CT was not destroyed by hydroxyl radical or hydroperoxide, but

• • it was destroyed by superoxide. CHP, which produces HO2 , HO2ˉ, and O2 ˉ, leads to a more efficient treatment because it provides oxidation, reduction, and desorption processes in one system (Furman et al., 2009). Many researchers have studied these oxygen species to understand their behavior during treatment systems.

Persulfate Chemistry

Persulfate is the newest application of an oxidant in ISCO projects. An inorganic compound,

2- persulfate dissolves in the aqueous phase to form peroxydisulfate anion (S2O8 ). It is capable of degrading a wide range of contaminants and has an oxidation-reduction potential of 2.01 V:

2- 2- S2O8 + 2eˉ → 2 SO4 (6)

The potential of oxidation-reduction for the persulfate anion is greater than the potential of oxidation-reduction for hydrogen peroxide (1.8 V) and permanganate (1.7 V), but it has a lower

potential than ozone (2.2 V).

Sodium persulfate (Na2S2O8) is the most popular persulfate salt in ISCO applications, due to its high solubility in water and low cost (e.g. compared to potassium persulfate). Because no harmful byproducts or residue are generated, sodium persulfate has a minimal environmental impact. The solubility of sodium persulfate is 73 g/100 ml @ 20°C, and it is denser than water.

These properties increase the mobility of a concentrated sodium persulfate solution, allowing it to easily move into the subsurface. Persulfate can degrade a vast array of contaminants and it is more stable than hydrogen peroxide and ozone (Huang et al., 2002), and it has a high longevity and can persist for weeks in the subsurface. Persulfate also can move through porous media and diffuse into low-permeability regions. In addition, Brown and Robinson (2004) found that persulfate does not react with soil organic matter (SOM) as quickly as permanganate.

Consequently, these features make persulfate one of the most used oxidants in remediation.

However, there are some limitations of persulfate. Sorbed contaminants are not desorbed and transformed by the persulfate even the persulfate is activated. Teel et al. (2009) found that enhanced desorption of sorbed contaminants was not achieved in the activated persulfate reactions. In addition, House (1962) noted that the persulfate anion does not react with recalcitrant compounds, such as TCE. House (1962) asserts that contaminant oxidation by persulfate is not achievable if the temperature is below 25°C. In this case, a catalyst is mandatory for oxidation to occur.

Persulfate activation potentially produces more oxidant and reductant oxygen species than unactivated persulfate. The most commonly identified species generated in persulfate activation

• • reactions include sulfate radical (SO4 ˉ), hydroperoxide anion (HO2ˉ), superoxide (O2 ˉ), and hydroxyl radical (OH•). When persulfate is catalyzed, sulfate radical is generated, which has a higher potential for oxidation (2.6 V) than that of ozone and slightly lower than that of hydroxyl radical (2.7 V). Sulfate radical is a strong oxidant that has a potential to degrade a limited number of organic compounds. The mechanism of degradation for sulfate radical is removing electrons from the organic compound and producing an organic radical cation. Huang et al.

(2002) demonstrated that the generation of sulfate radical in activated persulfate systems contributes to the generation of other strong species, such as hydroxyl radical (OH•), which results in a more powerful system that is able to degrade various organic contaminants. Sulfate radical is reactive with aromatic compounds possessing electron donating groups (e.g. hydroxyl

[-OH] and amino [-NH2]); however, aromatic compounds possessing electron withdrawing groups (e.g. chloro [-Cl] and nitro [-NO2]) are not oxidized by sulfate radical (Neta et al., 1977).

Several catalysts are used for persulfate activation. The most popular include heat, transition metals, and base activation. During heat activation of persulfate, sulfate radical is generated

(House, 1962):

2- • S2O8 + heat → 2 SO4 ˉ (7)

For persulfate heat activation, a minimum temperature of 35°C is required (House, 1962; Huang et al., 2002; Liang et al., 2003). Furthermore, Block et al. (2004) hypothesized that the minimum temperature requirement for thermal activation of persulfate varies from one target compound to another. For instance, they found that toluene degradation in thermally activated persulfate required a temperature of 20°C, while carbon tetrachloride and chlorobenzene degradation required 35°C and 45°C, respectively. Thermal activation of persulfate is not appropriate when

remediating large areas due to the high energy and installation costs of using large or multiple heating systems. However, it is suited to remediation of a small source area. In aqueous systems, but not soils, persulfate is usually heat activated. Liang et al. (2003) reported that oxidation of organic compounds by activated persulfate in soils requires greater temperature, persulfate volume, and time than in water.

Transition metals also have been used to activate persulfate. Ferrous iron is the most common metal catalyst used to activate persulfate; it is inexpensive, environmentally safe, and effective

(Zhao, 2010). Although persulfate can be activated using other transition metals such as copper and silver, these metals are toxic and cause environmental damage. Persulfate activation using iron (II) leads to sulfate radical generation and proceeds to persulfate reduction (Travina et al.,

1999):

2 +2 • 2 +3 S2O8 ˉ + Fe → SO4 ˉ + SO4 ˉ + Fe (8)

Although persulfate activation by iron (II) increases oxidation, excess iron causes a negative reaction between iron (II) and sulfate radical, as follows (Buxton et al., 1997):

• +2 2 +3 SO4 ˉ + Fe → SO4 ˉ + Fe (9)

As a result of the higher reaction rate (k= 4.6 x 109 M-1s-1) of iron (II) with sulfate radical, iron

(II) activation will be diminished and the oxidation performance will decline due to the rapid conversion of iron (II) to iron (III) and the corresponding rapid consumption of sulfate radical

(Liang et al., 2006). Therefore, iron chelate (Fe-ethylenediaminetetraacetic acid (EDTA)) or an optimum concentration of iron (II) is mandatory for persulfate activation (Liang et al., 2009).

Furthermore, iron (II) is involved in many chemical and physical reactions that decrease catalyst

transport and diffusion in the aquifer, limiting its subsurface mobility.

Base-activated persulfate is the most common activation method in ISCO remediation projects.

Furman et al. (2010) studied the mechanism of base activation of persulfate. They demonstrated that the hydrolysis of one persulfate molecule leads to the formation of hydroperoxide (HO2ˉ)

• (Reaction 10), which then reduces another persulfate molecule to form sulfate radical (SO4 ˉ) and

2ˉ sulfate anion (SO4 ). Hydroperoxide is simultaneously oxidized to superoxide radical anion

• (O2 ˉ) (Reaction 11).

2 2 + S2O8 ˉ + 2H2O → HO2ˉ + SO4 ˉ + 3H (10)

2 • 2 • + S2O8 ˉ + HO2ˉ → SO4 ˉ + SO4 ˉ + O2 ˉ + H (11)

Superoxide, which is generated in base-activated persulfate systems, contributes to the degradation of contaminants. Furman et al. (2010) used hexachloroethane as a superoxide probe and demonstrated that superoxide reactivity in base-activated persulfate reactions increases with increasing hydroperoxide generation. The popularity of this process has increased for ISCO remediation of soils and ground water because persulfate is more stable than hydrogen peroxide

(Furman et al., 2010). Hydroxyl radical is also generated by base activation of persulfate through the reaction of sulfate radical with hydroxide (Reaction 12).

• 2 • SO4 ˉ + OHˉ → SO4 ˉ + OH (12)

2ˉ • • Generating several reactive species (SO4ˉ, SO4 , O2 ˉ, OH , HO2ˉ) in base-activated persulfate reactions offers an advantage in degrading a vast array of contaminants in groundwater and soils.

However, base-activated persulfate does not produce enough superoxide for contaminant desorption (Teel et al., 2009).

Sodium Percarbonate Chemistry

Sodium percarbonate (PC) is a water-soluble (140 g/L) solid solution containing sodium carbonate and hydrogen peroxide with formula 2Na2CO3 · 3H2O2. It is commonly called solid hydrogen peroxide because it has a significant percentage of hydrogen peroxide (~30%). It is produced by the reaction between sodium carbonate and hydrogen peroxide. It has been used for different household cleaning and bleaching products such as OxiClean. When sodium percarbonate dissolves in water, it decomposes into sodium, carbonate, and hydrogen peroxide

(Reaction 13).

+ 2 2Na2CO3 · 3H2O2 → 4Na + 2CO3 ˉ + 3H2O2 (13)

Compared to liquid H2O2, sodium percarbonate is cheap, easy to transport because it is a low weight powdered solid, and safe because it is non-toxic to microbes and environmentally friendly. In aqueous systems, PC decomposes to peroxymonocarbonate (HCO4¯), which is highly reactive (Mekillop and Sanderson, 2005; Lin and Liu, 2008). HCO4¯ undergoes homolysis of the

• O−O bond, which generates several reactive species including carbonate radical (CO3 ¯),

• • perhydroxyl radical (HO2 ), and superoxide radical (O2 ˉ) (Reactions 14-16) (Lin and Liu, 2009):

• • HCO4¯ → CO3 ¯+ OH (14)

• • CO3 ¯+ H2O2 → HCO3¯ + HO2 (15)

• • + HO2 → O2 ¯ + H (16)

However, Merker et al. (2012) found that superoxide is a dominant reactive species in aqueous

PC systems. In addition, sodium percarbonate naturally provides an alkaline condition when it

dissolves in water due to the release of sodium carbonate, which increases the reactivity of superoxide (Furman et al., 2011). Furthermore, because it is effective over a wide pH range, sodium percarbonate degrades many organic compounds in water. Sodium percarbonate therefore has advantages for the treatment of chlorinated contaminants. De la Calle et. al. (2012) pointed out that sodium percarbonate is currently used as an oxidant method for ISCO.

The reactivity of sodium percarbonate with organic compounds has been widely studied. For example, Ando et al. (1986) reported that aniline and p-chloroaniline have been oxidized by sodium percarbonate to produce nitrobenzene and p-chloronitrobenzene. They also documented that styrene, cyclohexene, and cinnamyl alcohol are degraded by sodium percarbonate. In the degradation of tosylhydrazones to aldehydes and ketones, sodium percarbonate is highly reactive and the reaction rate is increased further with the addition of silica gel (Narayana et al., 1992).

Olah et al. (1991) found that benzene and toluene were oxidized by sodium percarbonate with trifluoromethanesulfonic acid. Solvents and acids are required in these systems, however. The reactivity of sodium percarbonate with aromatic compounds is dependent on the position of substituents in the aromatic ring. Kabalka et al., (1990) found that meta- and para-substituted benzonitrile are oxidized faster than ortho-substituted isomers. In summary, numerous organic contaminants (e.g. TCE, chlorinated aromatic hydrocarbons, polycyclic aromatic hydrocarbons

(PAHs), anthracene, phenanthrene) are degraded by sodium percarbonate (Yue-hua et al., 2011; de la Calle et. al., 2012; Cajal et al., 2012; Viisimaa, & Goi, 2014).

Objectives

The effectiveness of PC systems for generating a high flux of superoxide is well proved by

Merker et al. (2012). However, minimal studies have been published regarding the reactivity of superoxide generated in PC systems for remediation of halogenated contaminants and reduced compounds. In addition, the effect of the degree of halogenation and the halogen type on the rate of attack on halogenated compounds by PC-generated superoxide have not studied yet. The first objective of this research is:

 To survey a variety of contaminants that react with superoxide in PC systems (Chapter 2);

and

 To investigate the relationship between the degree of halogenation and the type of

halogen in the rate of attack on halogenated compounds by PC-generated superoxide

(Chapter 2).

Recent studies on superoxide reactivity reported that superoxide has the potential for enhanced desorption and destruction of sorbed contaminants in modified Fenton’s reactions having H2O2 concentration < 5 % (Smith et al., 2006; Corbin et al., 2007; Furman et al., 2009). Superoxide is a dominant reactive species in PC systems; and it is also generated in CHP and base-activated persulfate (BAP) systems. Additionally, minimal investigation of PC treatment of sorbed hydrophobic compounds has been accomplished to date. Therefore, the second objectives of the research is:

 To evaluate the potential for, and investigate trends in, the degradation of hydrophobic

contaminants sorbed onto silica gel in PC treatment systems (Chapter 3); and

 To compare PC, CHP, and BAP treatments (Chapter 3).

Superoxide reactivity increases alongside with increasing the surface area in aqueous superoxide

systems. Furman et al. (2009) documented that adding glass beads or the manganese oxide birnessite increased the reactivity of superoxide. In addition, adding finely crushed glass to bicarbonate-activated H2O2 systems led to increasing reactivity of superoxide (Xu et al., 2011).

Therefore, the presence of minerals or solids having a large surface areas in aqueous phases may increase the reactivity of superoxide and provide an effective and economical remediation for contaminants. The third objective of the research is:

 To investigate increasing reactivity of superoxide with solids and minerals (Chapter 4)

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CHAPTER 2

Survey of Contaminant Reactivity with Superoxide and Relationship Between Degree

of Halogenation and Type of Halogen in Rate of Attack by Superoxide

Introduction

Superoxide is a weak nucleophile and reductant. The reactivity and lifetime of superoxide in deionized water are much shorter than in organic solvents, because superoxide is strongly solvated by water molecules; and in deionized water, “it forms hydrogen bonds and undergoes rapid disproportionation that lowers its lifetime and reactivity” (Marklund, 1976; Bielski & Allen,

1977; Sawyer & Valentine, 1981; Afanas’ev, 1989). However, adding solvents that are less polar than water (e.g. hydrogen peroxide [H2O2]) to deionized water increases the reactivity of superoxide (Smith et al., 2004). It has been shown that superoxide is reactive with many organic compounds. Superoxide has the potential to degrade highly oxidized organics such as carbon tetrachloride (CT) (Smith et al., 2004; Teel & Watts, 2002). Nucleophilic species such as superoxide anion degrade some organic compounds by destroying electron poor moieties (e.g., carbonyl carbon) (David & Seiber, 1999). These species also degrade halogenated compounds such as CT and trichloroethylene (TCE). Furthermore, CT and other highly oxidized chlorinated organic compounds (e.g., hexachloroethane, bromotrichloromethane) are transformed by superoxide (Teel & Watts, 2002).

Superoxide is generated in several systems, such as catalyzed hydrogen peroxide propagations

(CHP), activated persulfate, bicarbonate-activated H2O2, and sodium percarbonate. Fenton’s reagent is a reaction between a dilute of H2O2 and a solution of metals or minerals that generates

hydroxyl radical in the reaction (OH•) (Haber & Weiss, 1934):

+2 • +3 H2O2 + Fe → OH + Fe (1)

CHP, or modified Fenton’s reagent, uses high concentrations of H2O2. As a result, several

• reactive oxygen species such as perhydroxyl radical (HO2 ), hydroperoxide anion (HO2ˉ), and

• superoxide radical anion (O2 ˉ) are produced (Reactions 2-5) (Watts et al., 1990; Barb et al.,

1951; Walling, 1975; De Laat and Gallard, 1999):

• • H2O2 + OH → HO2 + H2O2 (2)

+3 • + +2 H2O2 + Fe → HO2 + H + Fe (3)

• • + HO2 ↔ O2 ˉ + H pKa= 4.8 (4)

• +2 +3 HO2 + Fe → Fe + HO2ˉ (5)

Persulfate activation is also another method to produce superoxide. Base activation is the most popular and efficient pathway to activate persulfate for in situ chemical oxidation (ISCO) treatment processes. Base activation of persulfate under alkaline conditions was investigated and the mechanism elucidated by Furman et al. (2010). They reported that the base hydrolysis of a persulfate molecule contributes to the generation of hydroperoxide (HO2ˉ) (Reaction 6), and then

• another persulfate molecule is reduced by (HO2ˉ) forming sulfate radical (SO4 ˉ) and sulfate

2ˉ • anion (SO4 ). Then, superoxide radical anion (O2 ˉ) is generated as a result of oxidizing hydroperoxide (Reaction 7). Since persulfate is more stable than hydrogen peroxide, the use of this process has increased for ISCO remediation of soils and ground water (Furman et al., 2010).

2 2 + S2O8 ˉ + 2H2O → HO2ˉ + SO4 ˉ + 3H (6)

2 • 2 • + S2O8 ˉ + HO2ˉ → SO4 ˉ + SO4 ˉ + 2O2 ˉ + 3H (7)

Sodium percarbonate (PC) is another system generating superoxide. It is called solid H2O2 due to the solid powder containing a high percentage of H2O2. In aqueous PC systems, peroxymonocarbonate (HCO4¯) is generated as a result of decomposition of PC (McKillop and

Sanderson, 2005; Lin and Liu, 2008). Then, dissociation of the O−O bond of HCO4¯ contributes

• • • to producing carbonate radical (CO3 ¯), perhydroxyl radical (HO2 ), and superoxide radical (O2 ˉ)

(Lin and Liu, 2009). However, it has been shown that superoxide is the dominant reactive species in PC systems (Merker et al., 2012).

When superoxide is generated in PC reaction, it reacts with many organic compounds. For example, PC is capable of the degradation of aniline and p-chloroaniline to chloronitrobenzene and p-chloronitrobenzene (Ando et al., 1986). They also found that degradation of styrene, cyclohexene, and cinnamyl alcohol in PC systems. Additionally, PC treatment systems also have the potential for the degradation of organic compounds such as TCE, chlorinated aromatic hydrocarbons, polycyclic aromatic hydrocarbons (PAHs), anthracene, and phenanthrene (Yue- hua et al., 2011; de la Calle et. al., 2012; Cajal et al., 2012; Viisimaa, and Goi, 2014). The reactivity of sodium percarbonate with ethanes, ethenes, and chlorobenzenes has yet to be investigated. Therefore, the first objective of this work is to survey the reactivity of these contaminants with superoxide.

Degradation rates of halogenated compounds are related to the degree of halogenation and the type of halogen. Gillham and O’Hannesin (1994) observed that the rate of reaction of chlorinated compounds increases with increasing chlorination. Furthermore, nucleophilic reactions of halogenated compounds are related to the electronegativity of the halogen; higher

electronegativity contributes to a higher rate of degradation due to an attack by nucleophilic species such as OHˉ (Fukushima et al., 2011). The influence of halogen substitution on nucleophilic attack of halogenated compounds by superoxide generated in sodium percarbonate system will be the subject of the second objective of this work.

Methodology

Materials

Chloroform (CF), 1,2-dichloroethane (1,2-DCA), 1,2-dichloroethene (1,2-DCE), 1,1,1- trichloroethane (1,1,1-TCA), decane, dichloromethane, carbon tetrachloride (CT), pentachloroethane (PCA), 1,3,5-trinitrobenzene, trichloroethylene (TCE), tetrachloroethylene

(PCE), chlorobenzene, 1,2-dichlorobenzene, 1,2,4-trichlorobenzene, pentachlorobenzene, fluorotrichloromethane, bromotrichloromethane, percarbonate (PC), methyl ethyl ketone (MEK),

1,3-dinitrobenzene, 1,2,3,4-tetrachlorobenzene, and 1,4-dioxane were purchased from Sigma-

Aldrich (St. Louis, MO). Hexane, acetone, soluble potato starch, toluene, sulfuric acid, and nitrobenzene were purchased from J.T. Baker (Philipsburg, NJ). Methyl tert-butyl ether (MTBE), ammonium molybdate, and sodium thiosulfate were obtained from Fisher Scientific. Potassium iodide was obtained from Alfa Aesar (Ward Hill, MA).

Selection of Compounds

Several groups of contaminants were chosen for the investigation based on their expected

•¯ reactivity with superoxide (O2 ). The first group consisted of the methane derivatives dichloromethane, CF, and CT. The second group consisted of the ethane derivatives 1,2-DCA,

1,1,1-TCA, and PCA. The third group—1,2-DCE, TCE, and PCE—were ethene derivatives. The

fourth group, chlorobenzene derivatives, were chlorobenzene, 1,2-dichlorobenezne, 1,2,4- trichlorobenzene, 1,2,3,4-tetrachlorobenzene, and pentachlorobenzene. The fifth group, nitrobenzene derivatives, included 1,3-dinitrobenzene and 1,3,5-trinitrobenzene. Reduced compounds were also investigated to determine whether they were degraded in a superoxide- only system. This was group six and it contained 1,4-dioxane, MEK, MTBE, 1-pentanol, and toluene. The seventh group was selected to investigate the effect of halogen type (Br, Cl, F) on the reactivity of superoxide. This group contained fluorotrichloromethane, bromotrichloromethane, and carbon tetrachloride.

General Reaction Procedures

All degradation processes were conducted in triplicate in 20 mL volatile organic analysis (VOA) vials capped with polytetrafluoroethylene (PTFE) lined caps containing 0.5 M PC; control vials contained deionized water (DI) in place of the PC. For the methane derivatives, reaction vials contained 25 µM of contaminant. For the group of ethane derivatives, the vials contained 15 µM of the contaminants. Reaction vials of the ethene derivatives contained 10 µM of contaminant.

The chlorobenzene derivative reaction vials contained 2 µM of contaminant. The nitrobenzene derivative vials contained 0.1 mM of contaminant. The 1,4-dioxane, MEK, MTBE, 1-pentanol, and toluene reaction vials contained 5 mM of contaminant. The concentrations of each group were based on the solubility of the compounds and the instrument detective limit for quantification of each class of compound. A triplicate set of vials for the methane, ethane, ethene, chlorobenzene, and nitrobenzene compounds were extracted with hexane at each time point.

Decane was used to extract 1,4-dioxane, MEK, MTBE, 1-pentanol, and toluene in a triplicate set of vials at each time point.

Analysis

Hexane extracts containing chlorinated methanes, ethanes, ethenes, or chlorobenzenes were analyzed using a Hewlett-Packard 5890A gas chromatograph (GC) equipped with an electron capture detector (ECD) and with a 30 m × 0.53 mm capillary EQUITY-5 column. The chloromethane extracts had an injector temperature of 100 °C, the detector was 250 °C, the initial oven temperature was 50 °C, the program rate was 30 °C /min, and the final temperature was 170 °C. The chloroethane and chloroethylene extracts had an injector temperature of 220 °C, a detector temperature of 270 °C, an initial oven temperature of 100 °C, a program rate of 30 °C

/min, and a final temperature of 240 °C. For extracts containing chlorobenzene contaminants, the injector temperature was 300 °C, the detector temperature was 325 °C, the initial oven temperature was 60 °C, the program rate was 30 °C /min, and the final temperature was 300 °C.

Hexane extracts containing nitrobenzene contaminants and decane extracts containing 1,4- dioxane, MEK, MTBE, 1-pentanol, and toluene were analyzed using a gas chromatograph (GC) equipped with a flame ionization detector (FID) and fitted with 15 m × 0.53 mm capillary SPS-5 column. For nitrobenzene compounds, the injector temperature was 200 °C, the detector temperature was 250 °C, the initial oven temperature was 60 °C, the program rate was 30 °C

/min, and the final temperature was 180 °C. For 1,4-dioxane, the injector temperature was

240 °C, the detector temperature was 250 °C, the initial oven temperature was 60 °C, the program rate was 30 °C /min, and the final temperature was 240 °C. For MEK, MTBE, 1- pentanol, and toluene, the injector temperature was 220 °C, the detector temperature was 250 °C, the initial oven temperature was 60 °C, the program rate was 50 °C /min, and the final temperature was 240 °C. Microsoft Office Excel version 2013 was used to calculate the

averages, standard deviations, and standard errors for triplicate sets of reactions. KaleidaGraph version 4.0 was used to plot the figures with the standard errors.

Results and Discussion

Degradation of Chlorinated Contaminants

1) Degradation of Methane Derivatives

Degradation of 25 µM dichloromethane, CF, and CT in 0.5 M PC over 72 hours was studied

(Fig. 1). The loss of dichloromethane in the treatment and control systems was 58 % and 9 %, respectively. The loss of CF in the same systems with the same reaction time was 62 % (reaction system) and 6 % (control). Degradation of the final methane derivative, CT, over the course of the reaction was 71 % and 4 % in the treatment and control systems, respectively. These results demonstrate that dichloromethane, CF, and CT were degraded effectively in the PC treatment system in the aqueous phase. Merker et al. (2012) reported that superoxide is the dominant reactant generated in the PC system. Therefore, the findings prove that degradation of methane derivatives is accomplished by superoxide.

2) Degradation of Ethane Derivatives

Degradation of 15 µM 1,2-DCA, 1,1,1-TCA, and PCA in 0.5 M PC over 72 hours was investigated (Fig. 2). The loss of 1,2-DCA was 48 % in the treatment system and 6 % in the control system. The loss of 1,1,1-TCA over the same period in PC and control systems was 56 % and 4 %, respectively. The last derivative, PCA, was degraded in the PC system at 81 % and the loss in the control system was 4 %. The results illustrate that superoxide generated in the PC system is able to degrade ethane derivative contaminants in the aqueous phase.

3) Degradation of Ethylene Derivatives

Degradation of 10 µM 1,2-DCE, TCE, and PCE was achieved in the PC treatment system over

72 hours (Fig. 3). The degradation of 1,2-DCE was 51 % and 7 % in the reaction and control systems, respectively. The loss of TCE in the PC and control systems over the same period of time was 68 % and 6 %, respectively. The last ethylene compound, PCA, was degraded in the reaction and control systems at 87 % and 9 %, respectively. The results show that ethylene contaminants are degraded effectively by superoxide generated in PC systems.

4) Degradation of Chlorobenzene Derivatives

Degradation of 2 µM of chlorobenzene, 1,2-dichlorobenzene, 1,2,4-trichlorobenzene, 1,2,3,4- tetrachlorobenzene, and pentachlorobenzene in a 0.5 M PC system over 72 hours was investigated (Fig. 4). The loss of chlorobenzene in PC and control systems was 52 % and 10 %, respectively. Over the same time period, 1,2-dichlorobenzene was lost in the treatment system and in the control system at 57 % and 11 %, respectively. Degradation of 1,2,4-trichlorobenzene in the reaction and control systems was 74 % and 8 %, respectively. For 1,2,3,4- tetrachlorobenzene, the loss in the PC and control systems was 83 % and 9 %, respectively. The last compound in this group, pentachlorobenzene, was degraded in the treatment system at 85 %, while the loss in the control system was 11 %. The results document the capability of PC treatment system to generate superoxide to degrade chlorobenzene derivatives.

The results show that chlorinated organic compounds are degraded effectively in PC treatment systems, which generate superoxide. Destruction of halogenated carbon compounds by superoxide anion electrochemically in conventional aprotic solvents was investigated decades

ago (Sawyer and Roberts, 1983; Sawyer and Calderwood, 1984; Sugimoto et al., 1988; Kalu and

White, 1991; Sawyer et al., 1992). However, because aprotic solvents are volatile, using these applications for generation of superoxide was not practical for industrial applications (Hayyan et al., 2012). Teel and Watts (2002) found that the highly chlorinated compounds (e.g. CT, HCA, bromotrichloromethane) are transformed by superoxide in CHP systems. Therefore, using PC treatment systems for the degradation of highly halogenated contaminants is beneficial because it provides a high flux of superoxide, the responsible species for degrading these contaminants; also, PC systems are less expensive than CHP systems.

Effect of Degree of Halogenation

The effect of degree of halogenation on the rate of degradation in PC systems was investigated as shown in Figures 1-4. The rate of degradation for halomethanes was CT > CF > dichloromethane

(Fig. 1). The rate of degradation for chloroethanes was PCA > 1,1,1-TCA > 1,2-DCA (Fig. 2).

The rate of degradation for chloroethenes was PCE > TCE > 1,2-DCE (Fig. 3). The rate of degradation for chlorobenzenes was pentachlorobenzene > 1,2,3,4-tetrachlorobenzene > 1,2,4- trichlorobenzene > 1,2-dichlorobenzene > chlorobenzene (Fig. 4). The results show that the rate of degradation increases with increasing the degree of halogenation. These findings are in agreement with Gillham and O’Hannesin (1994) who observed that the rate of reaction of chlorinated compounds increases with increasing degree of chlorination. In addition, because superoxide is the only reactive species in PC system, the results illustrate that the rate of attack of superoxide increases with increasing degree of halogenation in the attacked compound.

Effect of Halogen Type on the Degradation Rate

The effect of halogen type on the rate of degradation in PC treatment systems was studied.

Figure 5 shows the degradation of 10 µM trichlorofluoromethane, CT, and bromotrichloromethane in 0.5 M PC treatment system. The loss of trichlorofluoromethane, CT, and bromotrichloromethane in the reaction systems was 90 %, 71 %, and 58 %, respectively, over 72 hours. The results demonstrate that the highest electronegativity compound (fluorine) was degraded more rapidly than the lower electronegativity compounds (chlorine, bromine).

Nucleophilic attack on halogenated compounds is related to the electronegativity of halogen; the higher electronegativity of a halogen contributes to a higher rate of degradation due to attack by nucleophilic species such as OHˉ (Fukushima et al., 2011). Therefore, due to generation of superoxide (nucleophilic species) in PC reactions, the highest electronegativity compound

(trichlorofluoromethane) was degraded more rapidly than the lower electronegativity compounds

(CT and bromotrichloromethane).

Hydrogen Peroxide (H2O2) Measurement

Sodium percarbonate contains a high amount of H2O2. The concentration of H2O2 in the PC systems was measured during the degradation processes for methane, ethane, ethylene, and chlorobenzene derivatives in 0.5 M PC. The average amount of H2O2 is presented in Figure 6. It shows that the % H2O2 increased and reached a peak of 0.79 % after 24 hours; then the % H2O2 decreased until reaching the lowest percentage at 72 hours. Because PC is used as a solid reagent, the percentage of H2O2 increases with increasing concentration of dissolved PC in the aqueous solution. After 24 hours, the peak of % H2O2 was reached when PC dissolved completely in the solution. After 24 hours, % H2O2 decreased due to the short longevity of %

H2O2 in aqueous phases (half life of days).

Degradation of Nitrobenzene Derivatives and Selected Reduced Compounds

Degradation of nitrobenzene derivatives (Fig. 7) and 1,4-dioxane, MEK, MTBE, 1-pentanol, and toluene (Fig. 8) in the PC treatment system was studied. The results show that nitrobenzene derivatives were not degraded in the PC system, which proves that superoxide does not have the ability to degrade nitrobenzene compounds (Fig. 7). Degradation of selected reduced compounds in PC system was also investigated (Fig. 8). The degradation of 1,4-dioxane was not achieved in the reaction systems. The loss of other reduced compounds, MEK, 1-pentanol, MTBE, and toluene, were 17%, 22%, 22%, and 24%, respectively, over 72 hours (Fig. 8). The obtained results demonstrate that MEK, MTBE and toluene were slightly removed in PC systems, however, the rate of loss was minimal. Therefore, PC system is not an appropriate choice for degrading of reduced compounds since it only generates superoxide, which does not react with reduced compounds.

Conclusion

Degradation of methane, ethane, ethylene, and chlorobenzene derivatives in 0.5 M PC was investigated. All chlorinated compounds were degraded by the PC system. Superoxide is the only significant reactive species in PC systems, and it is responsible for the degradation of chlorinated methanes, ethanes, ethylenes, and benzenes. In addition, it was found that the rate of degradation of halogenated compounds by superoxide generated in the PC system increased with greater degree of halogenation. Moreover, the highest electronegativity halogenated methane

(trichlorofluoromethane) was degraded more rapidly than the lower electronegativity

halogenated methanes (CT, bromotrichloromethane), which proves that the type of halogen affects the reactivity of superoxide.

The ability of PC systems to degrade nitrobenzene derivatives and other reduced compounds was also investigated. Our findings show that nitrobenzene derivatives and 1,4-dioxane were not degraded in the reaction systems, which shows that superoxide does not play a role in degradation of these contaminants in aqueous phases. The loss of MEK, MTBE, toluene, and 1- pentanol was insignificant in the reaction systems. Thus, PC system is not effective for degradation of reduced compounds in aqueous phases since the reactivity of superoxide with reduced compounds is negligible.

References

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Ando, T., Cork, D. G., & Kimura, T. (1986). “Sodium percarbonate” (SPC) as a hydrogen peroxide source for organic synthesis. Chemistry Letters, 15(5), 665-666.

Barb, W. G., Baxendale, J. H., George, P., & Hargrave, K. R. (1951). Reactions of ferrous and ferric ions with hydrogen peroxide. Transactions of the Faraday Society, 47, 462-499.

Bielski, B. H., & Allen, A. O. (1977). Mechanism of the disproportionation of superoxide radicals. The Journal of Physical Chemistry, 81(11), 1048-1050.

David, M. D., & Seiber, J. N. (1999). Accelerated hydrolysis of industrial organophosphates in water and soil using sodium perborate. Environmental Pollution, 105(1), 121-128.

Fukushima, M., Mizutani, Y., Maeno, S., Zhu, Q., Kuramitz, H., & Nagao, S. (2011). Influence of halogen substituents on the catalytic oxidation of 2, 4, 6-halogenated phenols by Fe (III)-tetrakis (p-hydroxyphenyl) porphyrins and potassium monopersulfate. Molecules, 17(1), 48-60.

Furman, O., Teel, A., & Watts, R. (2010). Mechanism of base activation persulfate. Environmental Science & Technology, 44, 6423-6428.

Gillham, R. W., & O'Hannesin, S. F. (1994). Enhanced degradation of halogenated aliphatics by zero‐valent iron. Groundwater, 32(6), 958-967.

Hayyan, M., Mjalli, F. S., Hashim, M. A., AlNashef, I. M., Al-Zahrani, S. M., & Chooi, K. L. (2012). Long term stability of superoxide ion in piperidinium, pyrrolidinium and phosphonium cations-based ionic liquids and its utilization in the destruction of chlorobenzenes. Journal of Electroanalytical Chemistry, 664, 26-32.

Kalu, E. E., & White, R. E. (1991). In situ degradation of polyhalogenated aromatic hydrocarbons by electrochemically generated superoxide ions. Journal of the Electrochemical Society, 138(12), 3656-3660.

Lin, J. M., & Liu, M. (2008). Chemiluminescence from the decomposition of

peroxymonocarbonate catalyzed by gold nanoparticles. The Journal of Physical Chemistry B, 112(26), 7850-7855.

Marklund, S. (1976). Spectrometric study of spontaneous disproportionation of superoxide anion radical and sensitive direct assay for superoxide dismutase. Journal of Biological Chemistry, 251, 7504–7507.

McKillop, A., Sanderson, W.R. (1995). Sodium perborate and sodium percarbonate: cheap, safe and versatile oxidising agents for organic synthesis. Tetrahedron, 51(22), 6145-6166.

Sawyer, D. T., & Valentine, J. S. (1981). How super is superoxide? Accounts of Chemical Research, 14(12), 393-400.

Sawyer, D. T., & Roberts Jr, J. L. (1983). U.S. Patent No. 4,410,402. Washington, DC: U.S. Patent and Trademark Office.

Sawyer, D. T., & Calderwood, T. S. (1984). U.S. Patent No. 4,468,297. Washington, DC: U.S. Patent and Trademark Office.

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Sugimoto, H., Matsumoto, S., & Sawyer, D. T. (1988). Degradation and dehalogenation of polychlorobiphenyls and halogenated aromatic molecules by superoxide ion and by electrolytic reduction. Environmental Science & Technology, 22(10), 1182-1186.

Teel, A., & Watts, R. (2002). Degradation of carbon tetrachloride by modified Fenton’s reagent. Journal of Hazardous Materials, B94, 179-189.

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Watts, R. J., Udell, M. D., Rauch, P. A., & Leung, S. W. (1990). Treatment of pentachlorophenol-contaminated soils using Fenton’s reagent. Hazardous Waste & Hazardous Materials, 7, 335-345.

Table 1. Groups of contaminants and their derivatives

Group Derivatives

Methane Dichloromethane, chloroform, and carbon tetrachloride (CT) 1,2-Dichloroethane, 1,1,1-trichloroethane (1,1,1-TCA), and Ethane pentachloroethane (PCA) Ethylene 1,2-Dichloroethylene, trichloroethylene (TCE), tetrachloroethylene (PCE) Chlorobenzene, 1,2-dichlorobenezne, 1,2,4-trichlorobenzene, 1,2,3,4- Chlorobenzene tetrachlorobenzene, and pentachlorobenzene Nitrobenzene Nitrobenzene (NB), 1,3-dinitrobenzene, and 1,3,5-trinitrobenzene Reduced 1,4-Dioxane, methyl ethyl ketone (MEK), methyl tetr-butyl ether Compound (MTBE), 1-pentanol, and toluene

List of Figures

Figure 1. Degradation of 25 µM dichloromethane, CF and CT in 0.5 M PC systems

Figure 2. Degradation of 15 µM 1,2-DCA, 1,1,1-TCA, and PCA in 0.5 M PC systems

Figure 3. Degradation of 10 µM 1,2-DCE, TCE, and PCE in 0.5 M PC systems

Figure 4. Degradation of 2 µM of chlorobenzene, 1,2-dichlorobenzene, 1,2,4- trichlorobenzene,1,2,3,4-tetrachlorobenzene, and pentachlorobenzene in 0.5 M PC systems

Figure 5. Degradation of 10 µM trichlorofluoromethane, CT, and bromotrichloromethane in 0.5 M PC systems

Figure 6. The average percentage of H2O2 during the degradation processes for methane, ethane, ethylene, and chlorobenzene derivatives in 0.5 M PC systems

Figure 7. Degradation of 2 mM nitrobenzene, 1,3-dinitrobenzene, and 1,3,5-trinitrobenzene in 0.5 M PC systems

Figure 8. Degradation of 5 mM 1,4-dioxane, MEK, 1-pentanol, MTBE and toluene in 0.5 M PC systems

Dichloromethane - control Dichloromethane CF - control CF CT - control CT

0.8

0.6

C/C 0.4

0.2

0 0 12 24 36 48 60 72 Time (hours)

Figure 1. Degradation of 25 µM CF, CT, and dichloromethane in 0.5 M PC systems

1,2-DCA - control 1,2-DCA 1,1,1-TCA - control 1,1,1-TCA PCA - control PCA

0.8

0.6

C/C 0.4

0.2

0 0 12 24 36 48 60 72 Time (hours)

Figure 2. Degradation of 15 µM 1,2-DCA, 1,1,1-TCA, and PCA in 0.5 M PC systems

1,2-DCE - control 1,2-DCE TCE - control TCE PCE - control PCE

0.8

0.6

C/C 0.4

0.2

0 0 12 24 36 48 60 72 Time (hours)

Figure 3. Degradation of 10 µM 1,2-DCE, TCE, and PCE in 0.5 M PC systems

Chlorobenzene - control Chlorobenzene 1,2-Dichlorobenzene - control 1,2-Dichlorobenzene 1,2,4-Trichlorobenzene - control 1,2,4-Trichlorobenzene 1,2,3,4-Tetrachlorobenzene - control 1,2,3,4-Tetrachlorobenzene Pentachlorobenzene - control Pentachlorobenzene 1

0.8

0.6

C/C 0.4

0.2

0 0 12 24 36 48 60 72 Time (hours)

Figure 4. Degradation of 2 µM of chlorobenzene, 1,2-dichlorobenzene, 1,2,4-trichlorobenzene, 1,2,3,4-tetrachlorobenzene, and pentachlorobenzene in 0.5 M PC systems

Bromotrichloromethane - control Bromotrichloromethane CT - control CT Trichlorofluoromethane - control Trichlorofluoromethane

0.8

0.6

C/C 0.4

0.2

0 0 12 24 36 48 60 72 Time (hours)

Figure 5. Degradation of 10 µM trichlorofluoromethane, CT, and bromotrichloromethane in 0.5 M PC systems

0.5M PC Control

0.8

0.6

2 0.4

% H

0.2

0 0 12 24 36 48 60 72 Time (hours)

Figure 6. The average percentage of H2O2 during the degradation processes for methane, ethane, ethylene, and chlorobenzenes derivatives in 0.5 M PC systems

Nitrobenzene - control 1,3-Dinitrobenzene - control 1,3,5-Trinitrobenzene - control Nitrobenzene 1,3-Dinitrobenzene 1,3,5-Trinitrobenzene 1

0.8

0.6

C/C 0.4

0.2

0 0 12 24 36 48 60 72 Time (hours)

Figure 7. Degradation of 2 mM nitrobenzene, 1,3-dinitrobenzene, and 1,3,5-trinitrobenzene in 0.5 M PC systems

1,4-Dioxane - control 1,4-Dioxane MEK - control MEK 1-Pentanol - control 1-Pentanol MTBE - control MTBE Toluene - control Toluene

0.8

0.6

C/C 0.4

0.2

0 0 12 24 36 48 60 72 Time (hours)

Figure 8. Degradation of 5 mM 1,4-dioxane, MEK, 1-pentanol, MTBE and toluene in 0.5 M PC systems

CHAPTER 3

Degradation of Hydrophobic Contaminants in Sodium Percarbonate Systems

Introduction

Soils and sediments contaminated with hydrophobic compounds have a significant impact on the environment. In addition to their environmental impacts, many of these contaminants are carcinogenic (Sullivan et al., 1992). Because the hydrophobicity of these contaminants is high

(water partition coefficients (log KOWs) > 4), they are sorbed onto solid phases. Therefore, the desorption rates are often very low (Quan et al., 2003). The transformation rates of sorbed organic compounds are limited by desorption rates. In hydrophobic compounds, desorption rates are generally slower than transformation rates. Sedlak and Andren (1994) demonstrated that the transformation of polychlorinated biphenyls (PCBs) is affected by sorption in the standard

Fenton’s reagent. In addition, the most popular oxidants (e.g. hydroxyl radical) cannot oxidize hydrophobic contaminants because these oxidants are not generated or not reactive in non- aqueous phases (Sheldon and Kochi, 1981; Sedlak and Andren, 1994). Furthermore, the biological degradation of sorbed contaminants is impossible due to the lack of bacterial reactivity in the solid phases and because the contaminants are protected from microbial attack by the sorption (Ogram et al., 1985; Aronstein et al., 1991). In situ chemical oxidation (ISCO) may play an important role for the desorption and destruction of hydrophobic sorbed contaminants because several ISCO methods generate superoxide, a species that has been shown to enhance the desorption and destruction of sorbed contaminants (Corbin et al., 2007).

The most efficient and popular ISCO systems are catalyzed H2O2 propagations (CHP) and

activated persulfate. CHP is a modification of the standard Fenton’s reaction. The standard

Fenton’s reagent is the reaction of dilute hydrogen peroxide (H2O2) with a solution of iron (II), which generates hydroxyl radical (OH•) (Haber and Weiss, 1934):

+2 • +3 H2O2 + Fe → OH + Fe (1)

Hydroxyl radical is a powerful oxidant that has the potential for the degradation of 98% of organic compounds in the aqueous phase (Haag and Yao, 1992). However, it is unreactive in non-aqueous phases (Sheldon and Kochi, 1981), so, it does not have the ability to degrade sorbed contaminants or non-aqueous phase liquids (NAPLs) (Sedlak and Andren, 1994; Watts et al.,

1999).

Several methods have been used to modify standard Fenton’s reactions to CHP. One such method is to increase the hydrogen peroxide concentration, as high as 1–5 M (Watts et al., 1990).

Modification of Fenton’s reactions are also applied by using some catalysts such as iron pyrophosphate (Wang & Brusseau, 1998), iron chelates (Pignatello & Baehr, 1994; Sun &

Pignatello, 1992), or oxyhydroxides (Miller & Valentine, 1994; Ravikumar & Gurol, 1994; Tyre et al., 1991; Watts et al., 1993, 1997). The use of concentrated hydrogen peroxide leads to

• propagation reactions, generating reactive oxygen species (e.g. perhydroxyl radical [HO2 ],

• hydroperoxide anion [HO2ˉ], and superoxide radical anion [O2 ˉ] (Watts et al., 1990; Barb et al.,

1951; Walling, 1975; De Laat and Gallard, 1999). The reactions are:

• • H2O2 + OH → HO2 + H2O2 (2)

+3 • + +2 H2O2 + Fe → HO2 + H + Fe (3)

• • + HO2 ↔ O2 ˉ + H pKa = 4.8 (4)

• +2 +3 HO2 + Fe → Fe + HO2ˉ (5)

Hydroperoxide anion is generated in the system due to oxidation of perhydroxyl radical

(Reaction 5). Hydroperoxide anion (HO2ˉ) is a strong nucleophile. Superoxide generation is

• • contributed by perhydroxyl radical (HO2 ) (Reaction 4). HO2 is a weak oxidant that reacts only with minimal organic compounds.

Catalyzed H2O2 propagations (CHP) is an efficient process for degrading hydrophobic contaminants. Watts et al. (1994) found that in CHP systems, when the H2O2 concentration increased to around 100 mM, the degradation of hexachlorobenzene was faster than its loss by gas-purge desorption, which is a measure of the maximum natural rate of desorption.

Furthermore, Watts et al. (1997) reported the desorption of a group of chlorobenzenes (e.g. 1,3,5- trichlorobenzene, 1,2,3,4-tetrachlorobenzene) sorbed onto hematite in CHP systems. In addition,

Watts et al., (1999) found sorbed hexachloroethane and hexachlorocyclopentadiene were degraded in a CHP system by non-hydroxyl radical species when the concentration of H2O2 was increased to more than 300 mM. Sorbed (polycyclic aromatic hydrocarbons) PAHs in soils and sediments are degraded by CHP using a high dosage of H2O2; the reaction conditions should be vigorous to achieve the optimum remediation and in soils or sediments having a high percentage of soil organic matter (oxidant scavenger) (Ferrarese et al., 2008). In addition, Corbin et al.

(2007) documented that superoxide is the responsible agent for enhanced desorption of dodecane in CHP system. These results demonstrate the potential for CHP to degrade hydrophobic contaminants and show that superoxide is the responsible for the enhanced desorption of sorbed chemicals.

Sodium percarbonate (PC) is a potentially important treatment option in the enhanced desorption of sorbed hydrophobic contaminants since it generates a significant amount of superoxide

(Merker et al., 2012). Several publications proved the effectiveness of PC treatment systems for remediation of chlorinated aromatic compounds and polycyclic aromatic hydrocarbons (e.g., acenaphthene, phenanthrene, anthracene, and fluoranthene) (de la Calle et. al., 2012; Cajal et al.,

2012; Viisimaa, and Goi, 2014). Minimal investigation of PC treatment of sorbed hydrophobic compounds has been accomplished to date. Therefore, the goal of this research is to evaluate the potential of PC for degradation of sorbed hydrophobic compounds (e.g. PCBs, insecticides, herbicides), and then to compare the PC treatment to CHP treatment and base-activated persulfate treatment.

Methodology

Materials

Hexane, acetone, sodium persulfate, sodium hydroxide, and iron (III)-ethylenediaminetetraacetic acid (EDTA) sodium salt (13% iron) were obtained from J.T. Baker (Philipsburg, NJ).

Hexachloropropene, 1,3-dichloropropene, atrazine, decane, hexachloro-1,3-butadiene, hexachlorocyclopentadiene, 1,2-dibromotetrachloroethane, lindane, hexachloroethane (HCA),

4,4ꞌ-dichlorodiphenyldichloroethane (4,4ꞌ-DDD), 4,4-dichlorodiphenyltrichloroethane, (4,4ꞌ-

DDT), dichlorodiphenyldichloroethylene (DDE), 2,4-dichlorophenoxyacetic acid (2,4-D), and

PC were obtained from Sigma-Aldrich (St. Louis, MO). 2,4,5 -Trichlorophenoxyacetic acid

(2,4,5-T) was obtained from CHEM-IMPEX International, Inc. (Wood Dale, IL). Bentazon,

2,2ꞌ,3-trichlorobiphenyl, and diuron were obtained from Ultra Scientific (North Kingstown, RI).

Silica gel 60, 0.062-0.105 mm (150-230 mesh) was obtained from Alfa Aesar (Ward Hill, MA).

2-Chlorobiphenyl, 2,3-dichlorobiphenyl, and 2,2ꞌ,3,3ꞌ-tetrachlorobiphenyl were obtained from

AccuStandard, Inc (New Haven, CT).

Selection of Hydrophobic Contaminants

Four groups of hydrophobic contaminants were selected for this research. Each group consisted of several of the most widespread hydrophobic contaminants with various degrees of chlorination, varied chemical structure, and varied KOW. The first group, organochlorine insecticides, included DDT, DDE, DDD, and lindane. The second group, chloroalkenes and chloroalkanes, contained hexachloropropene, 1,3-dichloropropene, hexachloroethane, 1,2- dibromotetrachloroethane, hexachloro-1,3-butadiene, and hexachlorocyclopentadiene. The third group was polychlorinated biphenyls (PCBs); it contained 2-chlorobiphenyl, 2,3- dichlorobiphenyl, 2,2ꞌ,3-trichlorobiphenyl, and 2,2ꞌ,3,3ꞌ-tetrachlorobiphenyl. The fourth group was herbicides; it contained atrazine, 2,4-D, 2,4,5-T, diuron, and bentazon.

Contaminant Loading

Because hydrophobic compounds are insoluble or only very slightly soluble in water, acetone was selected as the hydrophobic contaminant solvent. After making an acetone solution for each contaminant, 1 µmole of each solution was mixed with 100 g silica gel until the acetone evaporated completely, leaving the hydrophobic contaminant sorbed onto silica gel at 10

µmole/kg of silica gel.

Desorption Systems

Desorption of hydrophobic compounds from the silica gel was determined using the gas purge

method (GP). Gas purge desorption was applied in triplicate in 20 mL volatile organic analysis

(VOA) vials fitted with PTFE-lined caps. A Supelco ORBO-32 adsorbent tube was inserted to capture the contaminants. The reactors were prepared in the same manner as in the control reactions. The gas purge system was applied by supplying nitrogen purge gas into vials at a flow rate of 0.2 L min-1. At sampling times, the ORBO tubes were taken from the reactors and replaced instantly with new adsorbent tubes. The used tubes containing the volatilized contaminants were extracted with hexane or decane and then analyzed by gas chromatography.

Treatment Systems

Three systems were used to degrade the hydrophobic compounds, plus a control system:

Reactions with hydrophobic compounds and treatment systems were conducted in triplicate in 20 mL volatile organic analysis (VOA) vials with polytetrafluoroethylene (PTFE) lined caps. Each vial contained 1 g of silica gel loaded with the hydrophobic contaminant and 20 ml of reagent.

1) Catalyzed H2O2 Propagations (CHP)

The condition for CHP systems was 2 % H2O2 catalyzed using 1 mM iron (III) - EDTA.

2) Base-Activated Persulfate (BAP)

The condition for BAP systems was 0.5 M sodium persulfate activated by 1 M sodium hydroxide.

3) Sodium Percarbonate (PC)

The condition for PC systems was 0.5 M sodium percarbonate.

4) Control Samples

For control systems, each vial contained 1 g of silica gel loaded with hydrophobic contaminant and 20 ml of deionized water (DI) in place of reactive reagents (CHP, BAP, or PC).

Analysis

Silica gel containing the chloroalkene and chloroalkane compounds was extracted using hexane.

The extracts were then analyzed using a gas chromatograph (GC) equipped with an electron capture detector (ECD) and fitted with 30 m × 0.53 mm capillary EQUITY-5 column. The injector, detector, and initial oven temperatures were 220 °C, 220 °C, and 100 °C, respectively.

The program rate was 50 °C /min, and the final temperature was 160 °C.

The organochlorine insecticide, PCB, and herbicide compounds were extracted using decane.

The extracts were analyzed in GC equipped with a flame ionization detector (FID) and fitted with 15 m × 0.53 mm capillary SPS-5 column. The injector, detector, and initial oven temperatures were 240 °C, 250 °C, and 60 °C, respectively. The program rate was 30 °C /min, and the final temperature was 240 °C. Microsoft Office Excel 2013 was used to calculate the averages, standard deviations, and standard errors for triplicate sets of reactions. KaleidaGraph, v. 4.0, was used to plot the standard errors.

Results and Discussion

Degradation of Chloroalkanes and Chloroalkenes

The degradation of chloroalkane and chloroalkene contaminants in CHP, base-activated persulfate (BAP), and PC systems is shown in Figures 1-6. For the CHP system, 2 % H2O2 was

catalyzed using 1 mM iron (III) – EDTA. For the BAP system, 0.5 M sodium persulfate was activated by 1 M sodium hydroxide. For the PC system, 0.5 M sodium percarbonate was used for the experiment. Figure 1 shows that 1,2-dibromotetrachloroethane loss in PC, CHP, and BAP systems was 88 %, 83 %, and 56 % over 9 hours, respectively. The loss of 1,2- dibromotetrachloroethane in the GP and control systems was 16 % and 8 %, respectively. Figure

2 shows HCA loss in the PC, CHP, and BAP systems of 93 %, 86 %, and 49 %, respectively, over 9 hours. The loss of HCA in the GP and control systems was 16 % and 8 %, respectively.

These results demonstrate that sorbed 1,2-dibromotetrachloroethane and HCA are destroyed in the treatment systems used in this study. However, the rate of degradation in the PC system was slightly higher than in the CHP system and much higher than in the BAP system. Additionally, the results indicate that superoxide is responsible for degradation of the sorbed 1,2- dibromotetrachloroethane and HCA because the PC system generates only superoxide.

Moreover, the rate of HCA degradation was slightly faster than the rate of degradation of 1,2- dibromotetrachloroethane; this is likely because the electronegativity of chloride is higher than that of bromide. Because the rate of degradation of HCA and 1,2-dibromotetrachloroethane in

PC, CHP, and BAP systems was much higher than the rate of GP desorption, all were effective in the degradation process. Lastly, the results show that, under the conditions of this study, the PC system generates a higher flux of superoxide than the CHP or BAP systems.

1,3-Dichloropropene and hexachloropropene sorbed onto silica gel were also studied for their degradation in the treatment systems using the same reaction conditions. The loss of 1,3- dichloropropene in CHP, PC, and BAP treatment systems over 2 hours was 93 %, 61 %, and

19 %, respectively; the loss in GP and control systems was 18 % and 1 %, respectively (Fig. 3).

Figure 4 shows that the degradation of hexachloropropene in CHP, PC, and BAP systems was

93 %, 86 %, and 55 % over the same period of experiment, respectively, while the loss in GP and control systems was 14 % and 6 %.

The results show that 1,3-dichloropropene and hexachloropropene are effectively degraded in

CHP and PC systems. The fact that both compounds were degraded in the PC system suggests that superoxide can effectively degrade these sorbed hydrophobic contaminants. Moreover, the rates of removal of 1,3-dichloropropene and hexachloropropene were higher in the CHP system than in the PC system, indicating that the other reactive species produced in CHP may have minor roles in degradation of both compounds. In contrast, the rate of degradation of 1,3- dichloropropene in the BAP system (19 %) was almost the same as the rate of desorption in the

GP system (18 %); therefore, BAP is not effective in degrading sorbed 1,3-dichloropropene.

However, the BAP system has the potential for degradation of hexachloropropene. Lastly, because the degree of chlorination of hexachloropropene is higher than the degree of chlorination of 1,3- dichloropropene, the rate of degradation for hexachloropropene was higher than for 1,3- dichloropropene.

The degradation of hexachloro-1,3-butadiene, and hexachlorocyclopentadiene was investigated.

As shown in Figure 5, the loss of hexachloro-1,3-butadiene in CHP, PC, BAP, GP, and control systems in 2 hours was 61 %, 50 %, 6 %, 19 %, and 6 %, respectively. The results show that hexachloro-1,3-butadiene is degraded in CHP and PC systems. But the degradation rate in BAP system was lower than the rate of desorption in GP, which indicates that desorption of sorbed

hexachloro-1,3-butadiene did not occur in BAP reactions. The degradation rate of hexachlorocyclopentadiene in CHP, PC, BAP, GP, and the control systems for the course of the reaction was 92 %, 45 %, 16 %, 17 %, and 9 %, respectively (Fig. 6). The results also show that the ability of CHP and PC systems for the degradation of sorbed hexachlorocyclopentadiene. On the other hand, there was not degradation BAP reaction since the rate of degradation was lower than the rate of desorption in GP. These results are in agreement with Teel et al. (2009) who reported that hexachlorocyclopentadiene sorbed onto diatomaceous earth was not desorbed in the base-activated persulfate systems, however, it was desorbed and destructed effectively in CHP systems. The degradation rates of both compounds in the CHP system were markedly higher than the degradation rate in PC system.

Superoxide generated in CHP and PC systems significantly degraded sorbed hexachloro-1,3- butadiene and hexachlorocyclopentadiene in just 2 hours. Hydroxyl radical generated in the CHP

9 -1 -1 system may also play a role in destruction of hexachlorocyclopentadiene (kOH• = 2.6 x 10 M s )

(Haag and Yao, 1992) in the aqueous phase after its desorption by superoxide. The results also show that the rate of degradation of hexachloro-1,3-butadiene and hexachlorocyclopentadiene in the BAP system was lower than the rate of desorption in the GP system, demonstrating that BAP is inefficient at degrading these compounds. Therefore, using systems that produce a high flux of superoxide is recommended for degradation of sorbed hexachloro-1,3-butadiene and hexachlorocyclopentadiene.

Degradation of Organochlorine Insecticides

CHP, PC, and BAP systems were investigated for their ability to degrade four sorbed

organochlorine insecticides: DDE, DDD, lindane, and DDT. Figure 7 shows that DDE was degraded by BAP, CHP, and PC reactions at 44 %, 41 %, and 9 %, respectively, over 72 hours.

The loss of DDE in the GP and control systems was 18 % and 4 %, respectively. The loss of

DDD in BAP, CHP, PC, GP, and control systems was 64 %, 58 %, 16 %, 23 %, and 7 %, respectively, after 72 hours of reaction (Fig. 8). The percentage of lindane loss in BAP, CHP, and

PC, GP, and control systems was 48 %, 38 %, 10 %, 14 %, 8 %, respectively, over 72 hours (Fig.

9). The last organochlorine insecticide used was DDT. It was degraded in BAP, CHP, and PC treatment reactions at 45 %, 35 %, and 10 %, respectively. The loss of DDT in the GP and control systems was 19 % and 8 %, respectively (Fig. 10).

The results show that BAP and CHP are able to degrade sorbed organochlorine insecticides. In contrast, the PC treatment system did not degrade the organochlorine insecticides, because the rate of degradation of these insecticides in the PC system was lower than the rate of desorption in the GP system. Thus, superoxide did not degrade these contaminants. Moreover, the rate constant

8 -1 -1 for hydroxyl radical with lindane in the aqueous phase is kOH• = 4.2 x 10 M s (Haag and Yao,

1992), and hydroxyl radical is unreactive with sorbed contaminants. Destruction of sorbed lindane therefore follows desorption in the BAP and CHP systems that generate a vast array of reductants and oxidants in a single system.

Degradation of PCBs

A group of PCBs was treated using BAP, CHP, and PC treatments (Figs. 11-14). The loss of 2- chlorobiphenyl in BAP, CHP, PC, GP, and control systems was 48 %, 40 %, 2 %, 24 %, and 1 %, respectively, during 72 hours of reactions (Fig. 11); for 2,3-dichlorobiphenyl, it was 45 %, 32 %,

4 %, 18 %, and 2 %, respectively (Fig. 12); for 2,2’,3-trichlorobiphenyl, it was 44 %, 23 %, 2 %,

25 %, and 4 %, respectively (Fig. 13); and for 2,2’,3,3’-tetrachlorobiphenyl, it was 36 %, 24 %,

5 %, 18 %, and 3 %, respectively (Fig. 14). These results demonstrate that PCBs were not degraded in PC systems; therefore, superoxide did not react with PCBs.

The sorbed contaminants were degraded in BAP and CHP treatment systems, which show the ability of these systems to degrade PCBs. In contrast, the sorbed PCBs did not degrade in the PC system because their rate of degradation in the system was lower than the rate of desorption of sorbed PCBs in the GP system. Sato et al., (1993) found that hydroxyl radical has the potential to degrade PCBs adsorbed by sand, suggesting that hydroxyl radical was responsible for PCB degradation in the BAP and CHP systems. Furthermore, PCBs in the aqueous phase are degraded more rapidly than sorbed PCBs (Sedlak and Andren, 1994); therefore, the sorption negatively affected the remediation process by lowering the degradation rate.

Based on the obtained data, BAP degrades PCBs at a higher rate than CHP. Both systems produce similar reactive species except for sulfate radical, which is generated only in the BAP system. Sulfate radical and hydroxyl radical, therefore, may both be important in PCB degradation. The results are in agreement with Rastogi et al. (2009) who reported degradation of

PCBs in aqueous and non-aqueous phases by persulfate. To conclude, degradation of sorbed

PCBs is achievable in BAP and CHP systems due to the generation of several oxygen species in a single system. Superoxide does not degrade this group of contaminants.

Degradation of Herbicides

A group of herbicides (diuron, atrazine, bentazon, 2,4-dichlorophenoxyacetic acid (2,4-D), and

2,4,5-trichlorophenoxyacetic acid (2,4,5-T)) were studied to determine whether they are degraded by superoxide and to compare herbicide degradation in PC, CHP, and BAP systems.

The findings are shown in Figures 15-19. The loss of diuron in BAP, CHP, PC, GP, and control systems was 61 %, 42 %, 34 %, 19 %, and 4 % over 72 hours, respectively (Fig. 15). The results illustrate that the BAP and CHP systems effectively degrade sorbed diuron. The PC system was less efficient than the BAP and CHP systems; however, superoxide does degrade sorbed diuron to some degree in the PC system. In the literature, the second order rate constant for the reaction of diuron with hydroxyl radical in the aqueous phase is 4.6 x 109 M-1s-1 (De Laat et al., 1996).

Mazellier et al. (1997) reported that diuron is degraded efficiently by hydroxyl radical.

Therefore, sorbed diuron was thoroughly degraded in the BAP and CHP due to the combination of hydroxyl radical and superoxide anion in a single system.

The second herbicide chosen for this study was atrazine. Figure 16 shows that the removal rate of atrazine in BAP, CHP, PC, GP, and control systems was 44 %, 37 %, 5 %, 16 %, and 7 %, respectively, over 72 hours. The results demonstrate that BAP and CHP systems can degrade sorbed atrazine. However, atrazine did not react with superoxide generated in the PC system.

Numerous studies have reported that atrazine is effectively degraded by hydroxyl radical. The rate constant for the reaction of atrazine with hydroxyl radical in aqueous phase is 2.6 ± 0.4 x 109

M-1s-1 (Haag & Yao, 1992). Additionally, Guan et al. (2013) found that sulfate radical combined with hydroxyl radical are both important in degrading atrazine. Therefore, atrazine was degraded in the BAP system more rapidly than in the CHP system likely due to the generation of sulfate

radical.

Bentazon degradation was also studied. The loss of bentazon in CHP, BAP, PC, GP, and control systems was 66 %, 58 %, 28 %, 11 %, and 3 %, respectively, over 72 hours. CHP provided the greatest contaminant degradation, followed by BAP and then the PC system. The degradation of bentazon in the PC system demonstrates that superoxide is capable of degrading sorbed bentazon. However, the rate of degradation is not very high; it is possible that higher reagent concentrations may increase bentazon degradation in the PC system. Hydroxyl radical is likely responsible for bentazon degradation in the CHP and BAP systems; the rate constant for the reaction of bentazon with hydroxyl radical is 2.63 ± 0.1 x 109 M-1s-1 (Abdessalem et al., 2008).

However, superoxide generation in both systems is important to degradation since the loss of bentazon in PC system was 28 %, which illustrates that superoxide is a powerful agent for desorption and destruction of sorbed bentazon.

Degradation of 2,4-D, and 2,4,5-T was studied (Figs. 18-19). The percentage of loss of 2,4-D in

CHP, BAP, PC, GP, and the control was 58 %, 52 %, 47 %, 5 %, and 3 %, respectively, over 72 hours (Fig. 18). Over the same time of reaction, 2,4,5-T was degraded in CHP, BAP, PC, GP, and control systems at 46 %, 43 %, 28 %, 11 %, and 3 %, respectively (Fig. 19). The results show that the rate of degradation of 2,4-D is higher than that of 2,4,5-T. These findings are in agreement with Haag and Yao (1992) who reported a rate constant for the reaction of 2,4-D with hydroxyl radical of 5 x 109 M-1s-1 and a rate constant for the reaction of 2,4,5-T with hydroxyl radical of 4 x 109 M-1s-1.

The degradation of 2,4-D in the PC system was 47 %, which is slightly less than the degradation shown in the CHP and BAP systems, indicating that superoxide can degrade sorbed 2,4-D. The herbicide 2,4,5-T was also degraded in the PC system, but not as well as 2,4-D. In short, the rate of degradation is higher than the rate of desorption in a GP system, and, thus, CHP, PC, and BAP systems effectively degrade this group of herbicides. Atrazine, however, was only degraded in the CHP and BAP systems, demonstrating that superoxide does not degrade atrazine.

Conclusion Degradation of hydrophobic contaminants sorbed onto silica gel in PC, CHP, and BAP systems was investigated. The selected chloroalkene and chloroalkane compounds were effectively degraded using CHP and PC systems. Hexachloroethane and 1,2-dibromotetrachloroethane were the only compounds of this group degraded in the BAP system, where the rate of degradation was higher than the rate of desorption by the GP system. Organochlorine insecticides, the second group of hydrophobic compounds, were well degraded in the BAP and CHP systems. In contrast, the rate of degradation of insecticides in the PC system was lower than the rate of desorption in the GP system, showing that these compounds are not degraded by superoxide.

The third group of hydrophobic contaminants consisted of PCBs. All of the PCBs were degraded when treated in CHP and BAP reactions. However, the PCBs were not degraded in CHP and

BAP by superoxide; the rate of degradation of PCBs in the PC system was lower than the rate of desorption in the GP system. Herbicides were the final group examined in the present study.

Diuron, bentazon, 2,4-D, and 2,4,5-T were degraded effectively in CHP, BAP, and PC treatment systems with a degradation rate above the desorption rate of the GP system. However, atrazine was only degraded in the CHP and BAP systems, indicating that superoxide does not degrade

sorbed atrazine.

In conclusion, degradation of sorbed hydrophobic contaminants occurred with PC, CHP, and

BAP systems. However, the rate of degradation of each contaminant is dependent on the type of treatment. All of the selected sorbed hydrophobic contaminants were degraded effectively using the CHP system with 2 % H2O2 catalyzed by 1 mM iron (III) – EDTA. The PC and BAP systems degraded most contaminants. The BAP system, with 0.5 M sodium persulfate activated by 1 M sodium hydroxide, did not degrade 1,3-dichloropropene, hexachloropropene, hexachloro-1,3- butadiene, or hexachlorocyclopentadiene. The PC system, with 0.5 M sodium percarbonate, did not degrade PCBs, atrazine, or insecticides.

References

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Quan, H. N., Teel, A. L., & Watts, R. J. (2003). Effect of contaminant hydrophobicity on hydrogen peroxide dosage requirements in the Fenton-like treatment of soils. Journal of Hazardous Materials, 102(2), 277-289.

Rastogi, A., Al-Abed, S. R., & Dionysiou, D. D. (2009). Sulfate radical-based ferrous– peroxymonosulfate oxidative system for PCBs degradation in aqueous and sediment systems. Applied Catalysis B: Environmental, 85(3), 171-179.

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Sun, Y., & Pignatello, J. J. (1992). Chemical treatment of pesticide wastes. Evaluation of Fe (II) chelates for catalytic hydrogen peroxide oxidation of 2,4-D at circumneutral pH. Journal of Agricultural & Food Chemistry, 40, 322-327.

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Watts, R. J., Jones, A. P., Chen, P. H., & Kenny, A. (1997). Mineral catalyzed Fenton-like of sorbed chlorobenzene. Water Environment Research, 69, 269-275.

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Watts, R. J., Udell, M. D., Rauch, P. A., & Leung, S. W. (1990). Treatment of pentachlorophenol-contaminated soils using Fenton’s reagent. Hazardous Waste and Hazardous Materials, 7, 335-345.

Table 1. Selected groups of hydrophobic contaminants

Group Compounds Organochlorine DDT, DDE, DDD, and lindane Insecticides Hexachloropropene, 1,3-dichloropropene, hexachloroethane, 1,2- Chloroalkenes and dibromomtetrachloroethane, hexachloro-1,3-butadiene, and Chloroalkanes hexachlorocyclopentadiene Polychlorinated 2-Chlorobiphenyl, 2,3-dichlorobiphenyl, 2,2ˋ,3-trichlorobiphenyl, Biphenyls (PCBs) and 2,2ˋ,3,3ˋ-tetrachlorobiphenyl Herbicides Atrazine, 2,4-D, 2,4,5-T, diuron, and bentazon

List of Figures

Figure 1. Degradation of 1,2-dibromotetrachloroethane in CHP, BAP, and PC systems

Figure 2. Degradation of HCA in CHP, BAP, and PC systems

Figure 3. Degradation of 1,3-dichloropropene in CHP, BAP, and PC systems

Figure 4. Degradation of hexachloropropene in CHP, BAP, and PC systems

Figure 5. Degradation of hexachloro-1 3-butadiene in CHP, BAP, and PC systems

Figure 6. Degradation of hexachlorocyclopentadiene in CHP, BAP, and PC systems

Figure 7. Degradation of DDE in CHP, BAP, and PC systems

Figure 8. Degradation of DDD in CHP, BAP, and PC systems

Figure 9. Degradation of lindane in CHP, BAP, and PC systems

Figure 10. Degradation of DDT in CHP, BAP, and PC systems

Figure 11. Degradation of 2-chlorobiphenyl in CHP, BAP, and PC systems

Figure 12. Degradation of 2,3-dichlorobiphenyl in CHP, BAP, and PC systems

Figure 13. Degradation of 2,2',3-trichlorobiphenyl in CHP, BAP, and PC systems

Figure 14. Degradation of 2,2',3,3'-tetrachlorobiphenyl in CHP, BAP, and PC systems

Figure 15. Degradation of diuron in CHP, BAP, and PC systems

Figure 16. Degradation of atrazine in CHP, BAP, and PC systems

Figure 17. Degradation of bentazon in CHP, BAP, and PC systems

Figure 18. Degradation of 2,4-dichlorophenoxyacetic acid in CHP, BAP, and PC systems

Figure 19. Degradation of 2,4,5-trichlorophenoxyacetic acid in CHP, BAP, and PC systems

Control CHP GP PC BAP 1

0.8

0.6

C/C 0.4

0.2

0 3 Time (hours) 6 9

Figure 1. Degradation of 1,2-dibromotetrachloroethane in CHP, BAP, and PC systems

Control CHP GP PC BAP 1

0.8

0.6

C/C 0.4

0.2

0 0 3 6 9 Time (hours)

Figure 2. Degradation of HCA in CHP, BAP, and PC systems

Control PC GP CHP BAP 1

0.8

0.6

C/C 0.4

0.2

0 0 0.5 1 1.5 2 Time (hours)

Figure 3. Degradation of 1,3-dichloropropene in CHP, BAP, and PC systems

Control PC GP CHP BAP 1

0.8

0.6

C/C 0.4

0.2

0 0 0.5 1 1.5 2 Time (hours)

Figure 4. Degradation of hexachloropropene in CHP, BAP, and PC systems

Control PC GP CHP BAP 1

0.8

0.6

C/C 0.4

0.2

0 0 0.5 1 1.5 2 Time (hours)

Figure 5. Degradation of hexachloro-1 3-butadiene in CHP, BAP, and PC systems

Control PC GP CHP BAP

0.8

0.6

C/C 0.4

0.2

0 0 0.5 1 1.5 2 Time (hours)

Figure 6. Degradation of hexachlorocyclopentadiene in CHP, BAP, and PC systems

Control CHP GP BAP PC

0.8

0.6

C/C 0.4

0.2

0 0 12 24 36 48 60 72 Time (hours)

Figure 7. Degradation of DDE in CHP, BAP, and PC systems

Control CHP GP BAP PC

0.8

0.6

C/C 0.4

0.2

0 0 12 24 36 48 60 72 Time (hours)

Figure 8. Degradation of DDD in CHP, BAP, and PC systems

Control CHP GP BAP PC

0.8

0.6

C/C 0.4

0.2

0 0 12 24 36 48 60 72 Time (hours)

Figure 9. Degradation of lindane in CHP, BAP, and PC systems

Control CHP GP BAP PC 1

0.8

0.6

C/C 0.4

0.2

0 0 12 24 36 48 60 72 Time (hours)

Figure 10. Degradation of DDT in CHP, BAP, and PC systems

Control CHP GP BAP PC

0.8

0.6

C/C 0.4

0.2

0 0 12 24 36 48 60 72 Time (hours)

Figure 11. Degradation of 2-chlorobiphenyl in CHP, BAP, and PC systems

Control CHP GP BAP PC

0.8

0.6

C/C 0.4

0.2

0 0 12 24 36 48 60 72 Time (hours)

Figure 12. Degradation of 2,3-dichlorobiphenyl in CHP, BAP, and PC systems

Control CHP GP BAP PC

0.8

0.6

C/C 0.4

0.2

0 0 12 24 36 48 60 72 Time (hours)

Figure 13. Degradation of 2,2',3-trichlorobiphenyl in CHP, BAP, and PC systems

Control CHP GP BAP PC

0.8

0.6

C/C 0.4

0.2

0 0 12 24 36 48 60 72 Time (hours)

Figure 14. Degradation of 2,2',3,3'-tetrachlorobiphenyl in CHP, BAP, and PC systems

Control CHP GP BAP PC

0.8

0.6

C/C 0.4

0.2

0 0 12 24 36 48 60 72 Time (hours)

Figure 15. Degradation of diuron in CHP, BAP, and PC systems

Control CHP GP BAP PC 1

0.8

0.6

C/C 0.4

0.2

0 0 12 24 36 48 60 72 Time (hours)

Figure 16. Degradation of atrazine in CHP, BAP, and PC systems

Control BAP GP CHP PC

0.8

0.6

C/C 0.4

0.2

0 0 12 24 36 48 60 72 Time (hours)

Figure 17. Degradation of bentazon in CHP, BAP, and PC systems

Control BAP GP CHP PC

0.8

0.6

C/C 0.4

0.2

0 0 12 24 36 48 60 72 Time (hours)

Figure 18. Degradation of 2,4-dichlorophenoxyacetic acid in CHP, BAP, and PC systems

Control BAP GP CHP PC

0.8

0.6

C/C 0.4

0.2

0 0 12 24 36 48 60 72 Time (hours)

Figure 19. Degradation of 2,4,5-trichlorophenoxyacetic acid in CHP, BAP, and PC systems

CHAPTER 4

Increasing the Reactivity of Superoxide with Solids and Salts

Introduction

In situ chemical oxidation (ISCO) technologies are commonly used to remediate contaminated sites. Catalyzed hydrogen peroxide propagations (CHP), permanganate, and ozone are the most common ISCO technologies. CHP is an extension of the traditional Fenton’s reaction. In the traditional Fenton’s reaction, hydroxyl radical (OH•) is generated as a result of mixing dilute hydrogen peroxide (H2O2) with aqueous iron (II) at a pH of approximately 3 (Haber & Weiss,

1934):

+2 • +3 H2O2 + Fe → OH + Fe (1)

Hydroxyl radical is powerful oxidant that reacts with most organic compounds at or near diffusion controlled rates in the aqueous phase (Haag & Yao, 1992).

Fenton’s reagent is modified for CHP reactions by using high concentrations of hydrogen peroxide in the presence of soluble transition metals or minerals. As a result of this modification,

• • perhydroxyl radical (HO2 ), hydroperoxide anion (HO2ˉ), and superoxide radical anion (O2 ˉ) are produced in the system (Reactions 2-5) (Watts et al., 1990; Barb et al., 1951; Walling, 1975; De

Laat & Gallard, 1999).

• • H2O2 + OH → HO2 + H2O2 (2)

+3 • + +2 H2O2 + Fe → HO2 + H + Fe (3)

• • + HO2 ↔ O2 ˉ + H pKa= 4.8 (4)

• +2 +3 HO2 + Fe → Fe + HO2ˉ (5)

Superoxide is considered unreactive in deionized water because superoxide is strongly solvated by water molecules; it forms hydrogen bonds and undergoes rapid disproportionation (Marklund,

1976; Bielski & Allen, 1977; Sawyer & Valentine, 1981; Afanas’ev, 1989). However, superoxide is a strong nucleophile that becomes stable and reactive with the addition of solvents that are less polar than water (e.g. H2O2), which alter its solvation shell (Smith et al., 2004). In addition, adding other solvents that are less polar than water, such as acetone and ethylene glycol, in aqueous superoxide systems also increases the reactivity of superoxide. These findings indicate that the reactivity of superoxide in water increases with the addition of solvents of lower polarity than water.

The reactivity of superoxide in aqueous systems is also increased by adding solid surfaces to the system. Furman et al. (2009) found, for example, that adding birnessite with a surface area of

2 50.6 m /g to the H2O2 system increased superoxide reactivity. Similarly, adding glass beads to an aqueous superoxide system increased reactivity; the greater the surface area of the beads, the greater the reactivity. Xu et al. (2011) reported that superoxide reactivity was increased by adding finely crushing glass to increase the surface area in the system. These results demonstrate the relationship between superoxide reactivity and solids surface area. The objective of this research was to determine superoxide reactivity in response to the addition of several solids: iron oxides, manganese oxides, and trace minerals.

Methodology

Materials

Hexane, acetone, potassium permanganate (KMnO4), and hydrochloric acid (HCl) were obtained

from J.T. Baker (Philipsburg, NJ). Hexachloroethane and sodium percarbonate were obtained from Sigma-Aldrich (St. Louis, MO). Class V glass spheres of varying sizes were obtained from

MO-SCI Specialty Products, LLC (Rolla, MO). Birnessite (manganese oxide) was synthesized by the reduction of boiling potassium permanganate with concentrated hydrochloric acid

(McKenzie, 1989). The precipitate was washed repeatedly with deionized water until the pH exceeded 5. After the precipitate was washed, it was vacuum filtered and dried in an oven at

50 °C. The surface area of the birnessite was measured using Brunauer-Emmett-Teller (BET) analysis (Brunauer et al., 1938).

Bauxite, calcite, cuprite, hematite, ilmenite, manganite, pyrolusite, siderite, and willemite were purchased from D.J. Minerals (Butte, MT). They were received as 1 cm3 cubes and were crushed to a fine powder using a 150 mL capacity Spex shatter box with a grinder of hardened steel.

Mineral surface areas were measured using BET analysis under liquid nitrogen on a coulter SA

3100 (Carter et al., 1986). The surface areas of the minerals and their chemical formulas are listed in Table 1.

Probe Compound

-1 -1 Hexachloroethane (HCA) (kO2•ˉ = 400 M s ) (Afanas’ev, 1989) was used as a probe to detect the superoxide radical anion.

General Procedures

Percarbonate (PC) reactions were used to generate superoxide as the sole reactive species. The reactions containing HCA, PC, and various solids were conducted in triplicate in 20 ml volatile

organic analysis (VOA) vials with polytetrafluoroethylene (PTFE) lined caps. Each vial contained 16 ml of 2 μM HCA and 0.25 M PC, with or without the addition of solids. Glass bead reactions contained 3 g of varying sizes of glass beads. Mineral reactions contained 0.03 g birnessite, 0.05 g cuprite, calcite, and bauxite, or 0.1 g for the remainder of the minerals.

Controls were conducted in triplicate using deionized water (DI) in place of PC.

Analysis

HCA was extracted using hexane. The extracts were analyzed using a gas chromatograph (GC) equipped with an electron capture detector (ECD) and fitted with a 30 m × 0.53 mm capillary

EQUITY-5 column. The injector, detector, and initial oven temperatures were 220 °C, 220 °C, and 100 °C, respectively. The program rate was 50 °C /min and the final temperature was

160 °C. Microsoft Office Excel version 2013 was used to calculate the averages, standard deviations, and standard errors for the triplicate reaction sets. KaleidaGraph version 4.0 was used to plot the figures with standard errors.

Results and Discussion

Reactivity of Superoxide in PC Systems in the Presence of Glass Beads

Various sizes of glass beads were added to PC systems to investigate the effect of solid surfaces on the reactivity of superoxide. The reactivity of superoxide in a glass bead-PC system was quantified by the loss of HCA, a superoxide probe. Figure 1 shows that the degradation of HCA in the glass bead-PC system with 0.015, 0.03, and 0.059 m2/g glass beads surface area (SA) was higher than the degradation of HCA in the PC system with no solids. The greatest degradation rate was found when adding the finest glass beads with a surface area of 0.059 m2/g. These

results are in agreement with Furman et al., (2009) who investigated the effect of surface area on the reactivity of superoxide in the aqueous phase in a potassium superoxide system. Furman found that the reactivity of superoxide was proportional to the surface area of the glass beads.

Reactivity of Superoxide in PC Systems in the Presence of Iron Oxides

To determine the effect on superoxide reactivity of iron oxides, 0.1 g of goethite and hematite were added to PC systems. The reactivity of superoxide was quantified by HCA loss. Figure 2 shows that the greatest HCA degradation was achieved when adding goethite, which has a surface area of 1.4 m2/g. Adding hematite (SA = 0.3 m2/g) to the treatment system offered better

HCA degradation than in the PC system with no solids. The reactivity of superoxide in the hematite-PC system, however, was lower than the reactivity of superoxide in the goethite-PC system. Thus, superoxide reactivity was found to increase with greater surface area for iron oxides.

Reactivity of Superoxide in PC Systems in the Presence of Manganese Oxides

Birnessite, manganite, and pyrolusite were selected to determine superoxide reactivity relative to the surface area of manganese oxide minerals. As with the iron oxides, 0.1 g manganite and pyrolusite were used; however, birnessite was used at 0.03 g due to the higher SA of birnessite

(SA = 50.6 m2/g). Hexachloroethane loss was used to determine superoxide reactivity in each system. Superoxide reactivity with the addition of birnessite increased significantly and experimental vials of birnessite exploded rapidly. As shown in Figure 3, HCA was degraded more rapidly in the pyrolusite-PC system (5.2 m2/g surface area) than in the PC system, and was greatest with manganite (SA = 16 m2/g) and birnessite. In addition, 0.1 g manganite added to the

PC system increased superoxide reactivity at the same rate as 0.03 g of birnessite, which shows that the reactivity of superoxide increases with increasing SA. Thus, SA, rather than the nature of the mineral, is responsible for the increased reactivity in this case.

Reactivity of Superoxide in PC Systems in the Presence of Trace Minerals

Two groups of trace minerals were also studied for this research to determine their effect on superoxide reactivity. The first group consisted of siderite, willemite, and ilmenite; 0.1 g each of these minerals was added to PC systems. The greatest degradation of HCA was accomplished with the addition of siderite (SA = 6.8 m2/g) (Fig. 4). Hexachloroethane degradation with added willemite (SA = 1.8 m2/g) or ilmenite (SA = 1.7 m2/g) was greater than in a PC-only system, but lower than in the siderite-PC system. Furthermore, the reactivity of superoxide in both willemite-

PC and ilmenite-PC systems was similar; because both provide the same surface area in the systems, these results indicate that the increased reactivity in aqueous systems is proportional to the surface area added to the system.

The second group of trace minerals consisted of cuprite, calcite, and bauxite. Due to their high surface area, 0.05 g of these minerals was added to PC systems. Reactivity was quantified by

HCA loss from the system. As shown in Figure 5, the greatest loss of HCA was achieved with cuprite, followed by calcite and then bauxite. Because cuprite also provides the largest surface area (49.5 m2/g), followed by calcite (38 m2/g) and then bauxite (28 m2/g), the results indicate that superoxide reactivity in water increases as a function of the surface area of minerals; the nature of the mineral has no role in increasing of superoxide reactivity in mineral-PC systems.

The Relationship Between Surface Area and HCA Degradation Rate Constant in PC Systems

The rate constants for HCA degradation in glass bead and mineral PC systems were measured to further document the increased reactivity of superoxide as a function of the surface area of glass beads and minerals. The mass of added solids and minerals and their SAs in each vial are listed in Table 2. The relationship between the rate constants for HCA degradation in glass bead and mineral PC systems (y-axis) and the total SAs of solids and minerals in the systems

(x-axis) are shown in Figure 6. The rate constant increased with increased surface area of the added minerals or solids in the system. The rate constant results were as follows: cuprite > calcite > manganite > birnessite > bauxite > siderite > pyrolusite > GB (0.059 m2/g) = willemite = ilmenite < goethite < GB (0.030 m2/g). The lowest rate constants were found in hematite and glass beads (0.015 m2/g) which had the smallest surface areas of the solids and minerals studied.

The results illustrate that the SAs of the solids and minerals, rather than the nature of the solids and minerals, are responsible for the observed increase in superoxide reactivity in glass bead-

PC and mineral-PC systems. According to Furman et al. (2009), “the possible mechanisms for increasing the reactivity of superoxide with increasing surface areas include increased life-time of superoxide in the solid surface microenvironment, enhanced juxtaposition of organic compounds (e.g. HCA) with superoxide at solid surfaces, catalytic effects at solid surfaces, and changes in the solvation shell morphology of sorbed superoxide.”

Conclusion

The present research demonstrates increased superoxide reactivity in aqueous superoxide systems with the addition of solids or minerals due to increased SA. The results indicate that increased surface area—not the nature of the solid or mineral—is responsible for the increase in superoxide reactivity in aqueous PC systems. The linear relationship between the rate constants for HCA degradation and the actual solid or mineral surface areas demonstrates that providing surfaces for superoxide enhances its reactivity in the system. The largest rate constants for HCA degradation were found with minerals with large surface areas, such as cuprite, calcite, birnessite, manganite, and bauxite. On the other hand, using minerals or solids with small surface areas also increased the reactivity of superoxide in PC systems, but to a lesser degree. Therefore, the presence of minerals or any other solid in a superoxide-system increases treatment effectiveness in contaminated sites.

References

Afanas’ev, I. B. (1989). Superoxide ion: Chemistry and biological implication. (Vol. 1). Boca Raton, FL: CRC Press.

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Bielski, B. H., & Allen, A. O. (1977). Mechanism of the disproportionation of superoxide radicals. The Journal of Physical Chemistry, 81(11), 1048-1050.

Brunauer, S., Emmett, P. H., & Teller, E. (1938). Adsorption of gases in multimolecular layers. Journal of the American Chemical Society, 60(2), 309-319.

Carter, D. L., Mortland, M. M., & Kemper, W. D. (1986). Specific surface. Soil Science Society of America, 413-423.

Haag, W. R., & Yao, C. C. D. (1992). Rate constants for reaction of hydroxyl radicals with several drinking water contaminants. Environmental Science & Technology, 26(6), 1005- 1013.

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Sawyer, D. T., & Valentine, J. S. (1981). How super is superoxide? Accounts of Chemical Research, 14(12), 393-400.

Watts, R. J., Udell, M. D., Rauch, P. A., & Leung, S. W. (1990). Treatment of pentachlorophenol-contaminated soils using Fenton’s reagent. Hazardous Waste & Hazardous Materials, 7, 335-345. 95

Walling, C. (1975). Fenton’s reagent revisited. Accounts of Chemical Research, 8(4), 125-131.

Xu, A., Li, X., Hui, X., & Yin, G. (2011). Efficient degradation of organic pollutants in aqueous solution with bicarbonate-activated hydrogen peroxide. Chemosphere, 82, 1190- 1195.

Table 1. Minerals formulas and surface areas.

Minerals Formula Surface area (m2/g)

Birnessite MnO2 50.6 Bauxite Al(OH)3 28.8 Calcite CaCO3 38.0 Cuprite Cu2O 49.5 Goethite FeO(OH) 1.4 Hematite Fe2O3 0.3 Ilmenite FeTiO3 1.7 Manganite MnO(OH) 16 Pyrolusite MnO2 5.2 Siderite FeCO3 6.8 Willemite Zn2SiO4 1.8

Table 2. Mass (g) and surface area (m2) of added solids and minerals

Solid or Mineral Mass of added amount (g) Surface area added (m2) Glass Bead (80-100 mesh) 3 0.045 Glass Bead (170-200 mesh) 3 0.09 Glass Bead (325-400 mesh) 3 0.17 Birnessite 0.03 1.518 Bauxite 0.05 1.45 Calcite 0.05 1.9 Cuprite 0.05 2.475 Goethite 0.1 0.14 Hematite 0.1 0.03 Ilmenite 0.1 0.17 Manganite 0.1 1.6 Pyrolusite 0.1 0.52 Siderite 0.1 0.68 Willemite 0.1 0.18

List of Figures

Figure 1. Effect of adding 3 g glass beads (GB) on degradation of 2 μM HCA in 0.25 M PC systems

Figure 2. Effect of adding 0.1 g iron minerals on degradation of 2 μM HCA in 0.25 M PC systems

Figure 3. Effect of adding 0.1 g pyrolusite and manganite and 0.03 g birnessite on degradation of 2 μM HCA in 0.25 M PC systems

Figure 4. Effects of adding 0.1 g trace minerals on degradation of 2 μM HCA in 0.25 M PC systems

Figure 5. Effects of adding 0.05 g trace minerals on degradation of 2 μM HCA in 0.25 M PC systems

Figure 6. Relationships between the actual surface areas of glass beads (GB) and minerals and rate constants for degradation of 2 μM HCA in 0.25 M PC systems

Control GB (200) PC GB (300) GB (100)

0.8

0.6

C/C 0.4

0.2

0 0 1 2 3 4 5 6 Time (hours)

Figure 1. Effect of adding 3 g glass beads (GB) on degradation of 2 μM HCA in 0.25 M PC systems

Control Hematite PC Goethite

0.8

0.6

C/C 0.4

0.2

0 0 1 2 3 4 5 6 Time (hours)

Figure 2. Effect of adding 0.1 g iron minerals on degradation of 2 μM HCA in 0.25 M PC systems

100

Control Manganite PC Birnessite Pyrolusite 1

0.8

0.6

C/C 0.4

0.2

0 0 1 2 3 4 5 6 Time (hours)

Figure 3. Effect of adding 0.1 g pyrolusite and manganite and 0.03 g birnessite on degradation of 2 μM HCA in 0.25 M PC systems

101

Control Willemite PC Siderite Ilmenite 1

0.8

0.6

C/C 0.4

0.2

0 0 1 2 3 4 5 6 Time (hours)

Figure 4. Effects of adding 0.1 g trace minerals on degradation of 2 μM HCA in 0.25 M PC systems

102

Control Calcite PC Cuprite Bauxite 1

0.8

0.6

C/C 0.4

0.2

0 0 1 2 3 4 5 6 Time (hours)

Figure 5. Effects of adding 0.05 g trace minerals on degradation of 2 μM HCA in 0.25 M PC systems

103

0.35

y = 0.045214 + 0.10348x R= 0.99273 Cuprite 0.3

0.25 Calcite

HCA

K 0.2 Birnessite Manganite Bauxite

0.15

Siderite

Pyrolusite 2 0.1 4 1: GB (0.015 m /g) 3 2: Hematite 1 2 GB (0.059 m /g) 2 Willemite 3: GB (0.030 m /g) Ilmenite 4: Goethite 0.05 2

0 0.5 1 1.5 2 2.5 Actual Surface Area (m2)

Figure 6. Relationships between the actual surface areas of glass beads (GB) and minerals and rate constants for degradation of 2 μM HCA in 0.25 M PC systems

104