Treating Metals in Acid Mine Drainage Using Slow-Release Hydrogen Peroxide

Home , Acid mine drainage

A thesis presented to

the faculty of

the College of Arts and Sciences of Ohio University

In partial fulfillment

of the requirements for the degree

Master of Science

Samuel A. Miller

August 2015

2 This thesis titled

Treating Metals in Acid Mine Drainage Using Slow-Release Hydrogen Peroxide

SAMUEL A. MILLER

has been approved for

the Department of Geological Sciences

and the College of Arts and Sciences by

Eung Seok Lee

Associate Professor of Geological Sciences

Robert Frank

Dean, College of Arts and Sciences 3 ABSTRACT

MILLER, SAMUEL A., M.S., August 2015, Geological Sciences

Treating Metals in Acid Mine Drainage Using Slow-Release Hydrogen Peroxide

Director of Thesis: Eung Seok Lee

Metal concentrations from acid mine drainage (AMD) pose a significant threat to aquatic systems worldwide as a result of past and current mining operations. This study tested the efficacy of using slow-release hydrogen peroxide (SR-HP) to oxidize and remove ferrous iron (Fe2+) from AMD. Fenton’s reagent forms from a mixture of

2+ hydrogen peroxide (H2O2) and Fe available from AMD, creating an advanced oxidation process. Twenty-eight SR-HP forms were developed by dispersing sodium percarbonate

(Na2CO3 1.5H2O2) salts in a polymeric matrix. The SR-HP forms released H2O2 in flowing water at a peak release rate of 0.05 – 52.1 mg min-1 during the initial hour and

-1 continued to release H2O2 at a lower, stable release rate (0.02 – 1.5 mg min ) from a period of days to weeks depending on salt : binding agent mixing ratios. Oxidant : resin mixing ratios in addition to surface area were primary factors impacting the release profiles from the laboratory column leaching tests. Proof-of-concept iron removal tests indicate that SR-HP forms can efficiently remove Fe2+ from AMD within one minute.

+2 2+ Ideal [Fe ]/[H2O2] ratios for >80% Fe removal clustered around 2, with decreasing

Fe2+ removal as the ratio increases. A small-scale field test demonstrated the efficacy of

SR-HP at oxidizing Fe2+. Ferrous iron concentrations were reduced by 80% within the first hour of treatment. These results suggest feasibility of using SR-HP to treat oxidizable metals in AMD water. Further development of SR-HP forms with higher 4 release rates, longer durations, stronger binder, and improved reproducibility could be possible.

5 DEDICATION

To my parents and grandparents

6 ACKNOWLEDGEMENTS

I would like to thank The Korean Institute of Geoscience and Mineral Resources

(KIGAM), Ohio Department of Natural Resources – Division of Mineral Resources, and the Monday Creek Restoration Project for funding this project, sample analysis, and aiding in selection of potential proof-of-concept tests. I would like to thank my advisor

Dr. Eung Seok Lee for giving me this opportunity and providing me with insight along the way. I would also like to thank fellow graduate students who have been willing to listen to my concerns and offer thoughtful advice.

7 TABLE OF CONTENTS

Page

Abstract ………………………………………………………………………………...…3

Acknowledgements ……………………………………………………………………….6

List of Tables……………………………………………………………………………...9

List of Figures……………………………………………………………………………10

Chapter 1: Introduction…………………………………………………………………..13

1.1 Formation of AMD…………………………………………..…………………..16

1.2 Advanced Oxidation Reactions…………………………………………………..18

1.2.1 Fenton’s Reagent…………………………………………………………..19

1.3 Slow Release Systems……………………………………………………………20

1.4 Study Site………………………………………………………………………...22

1.5 Study Objectives…………………………………………………………………26

Chapter 2: Materials and Methods……………………………………………………….28

2.1 Material…………………………………………………………………………..28

2.2 Field Measurements……………………………………………………………...28

2.3 Designing Slow Release Forms………………………………………………….29

2.4 Estimating Release Rate: Column Leaching Tests………………………………30

2.4.1 Determination of Hydrogen Peroxide Concentration ……………………..32

2.4.2 Determination of Ferrous Iron Concentration ……………………………..33

2.5 Proof-of-Concept Ferrous Iron Removal Tests…………………………………..35

2.6 Small-scale Field Application Test………………………………………………35

Chapter 3: Results and Discussion……………………………………………………….37

3.1 Study Site………………………………………………………………………...37 8 3.1.1 Baseline Chemical Sampling………………………………………………38

3.2 Characterizing Slow Release Systems…………………………………….……..40

3.2.1 Slow Release Form Dimensions…………………………………………...40

3.2.2 Column Release Tests……………………………………………………...41

3.2.3 Oxidant Release Efficiency / Recovery Rate………………………………62

3.3 Proof-of-Concept Iron Removal Tests…………………………………………...64

3.4 Small-scale Field Application Test………………………………………………67

Chapter 4: Conclusion…………………………………………………………………....84

References………………………………………………………………………………..86

Appendix A: H2O2 fluxes from 6 SR-HP column tests over two weeks…………………90

+2 Appendix B: Weight of sodium percarbonate (kg/day) necessary for a [Fe ]/[H2O2] 2:1 for different Fe2+ concentrations (mg L-1) and discharges (L s-1)……………………..…91

9 LIST OF TABLES

Page

Table 1-1. Redox potential of common oxidants ………………………………………..18 Table 1-2. Historic water chemistry and discharge measurements at study site BH00690…………………….....…………………………………….…………………..26 Table 3-1. Historic discharge and chemical parameter measurements at BH00690 (http://watersheddata.com) ……………………..………..…………..…….…………….38 Table 3-2. Aqueous chemistry data for BH006900 sampled on 6/10/2014….……….….39 Table 3-3. Properties of slow-release forms made with sodium percarbonate and casting resin……………………………….….……………………………………………..……40 Table 3-4. P-values from two-sided T-tests performed between first order decay constants between SR-HP 23-28…………………….…….……………………………….……….62 Table 3-5: Properties of SR-HP forms used in field removal demonstration…….……...65 Table 3-6: Percentage of dissolved oxidant from SR-HP forms removed from field demonstration……………………………………………………………………….……69 Table 3-7: Discharge measured during the small-scale field demonstration at BH00690 and upstream and downstream measurements…………………………………..……….69

10 LIST OF FIGURES

Page

Figure 1-1. Locations of the three Appalachian Basin coal regions (from USGS, 2002).14 Figure 1-2. Slow release diffusion modeled in cross sectional view. (Lee and Schwartz, 2007a)..………………… ……………………………………………………………….21 Figure 1-3. Location of study site BH00690 within the Monday Creek Watershed (adapted fromMonday Creek Reclamation Project)……………………………………..23 Figure 1-4. Location of study site BH00690 in relation to Jobs New Pittsburg Rd and Brush Fork,and other AMD sampling points stored on NPS website (http://watersheddata.com)……………………………………………………………….24 Figure 1-5. Solubility of major and minor metals at different pHs (modified from Wei et al., 2005)…………………………………………………………………………………25 Figure 2-1: Release kinetics of SR-HP forms at 8 mL min -1 column test (adapted from Tong, 2013) …………………………………………………………………….………..31 Figure 2-2. Standard curve for Hydrogen Peroxide concentration determination……….33 Figure 2-3. Standard curve for Ferrous Iron concentration determination………………34 Figure 3-1: Small-scale field demonstration site, BH006900, located within the Monday Creek watershed. Baseline chemical sampling on June 10th, 2014…………………….37 Figure 3-2a. H2O2 release profile from a 1.66 : 1 SR-1 form during a column leaching test using a 7 ml min-1 flow rate………………………………………………………...42 Figure 3-2b. H2O2 release profile from a 2.5 : 1 SR-2 form during a column leaching test using a 7 ml min-1 flow rate…………………………………………………………...... 43 Figure 3-2c. H2O2 release profile from a 3.33 : 1 SR-3 form during a column leaching test using a 7 ml min-1 flow rate. …………………………………………….…………43 Figure 3-3a. H2O2 release profile from a 5 : 1 SR-4 form during a column leaching test using a 7 ml min-1 flow rate. …………………………………………………..………..44 Figure 3-3b. H2O2 release profile from a 5 : 1 SR-5 form during a column leaching test using a 7 ml min-1 flow rate. ……………………………………………………………45 Figure 3-3c. H2O2 release profile from a 6 : 1 SR-6 form during a column leaching test using a 7 ml min-1 flow rate. ……………………………………………………………45 Figure 3-3d. H2O2 release profile from a 6 : 1 SR-7 form during a column leaching test using a 7 ml min-1 flow rate. ……………………………………………………………46 Figure 3-3e. H2O2 release profile from a 5 : 1 SR-8 form during a column leaching test using a 7 ml min-1 flow rate. ……………………………………………………………46 Figure 3-3f. H2O2 release profile from a 5 : 1 SR-9 form during a column leaching test using a 7 ml min-1 flow rate. ………………………………………………...………….47 Figure 3-3g. H2O2 release profile from a 5 : 1 SR-10 form during a column leaching test using a 7 ml min-1 flow rate. ……………………………………………………………47 Figure 3-4a. H2O2 release profile from a 4.2 : 1 SR-11 form during a column leaching test using a 7 ml min-1 flow rate. ……………………………………………………….50

11 Figure 3-4b. H2O2 release profile from a 3.8 : 1 SR-12 form during a column leaching test using a 7 ml min-1 flow rate. ………………………………….……………………51 Figure 3-4c. H2O2 release profile from a 3.5 : 1 SR-13 form during a column leaching test using a 7 ml min-1 flow rate. …………………………………….…………………51 Figure 3-4d. H2O2 release profile from a 3.7 : 1 SR-14 form during a column leaching test using a 7 ml min-1 flow rate. ……………………………………………………….52 Figure 3-4e. H2O2 release profile from a 3.5 : 1 SR-15 form during a column leaching test using a 7 ml min-1 flow rate. ……………………………………….………………52 Figure 3-4f. H2O2 release profile from a 3.5 : 1 SR-16 form during a column leaching test using a 7 ml min-1 flow rate. ……………………………………………………….53 Figure 3-4g. H2O2 release profile from a 3.75 : 1 SR-17 form during a column leaching test using a 7 ml min-1 flow rate. ……………………………………………………….53 Figure 3-5a. H2O2 release profile from a 3.75 : 1 SR-18 form during a column leaching test using a 7 ml min-1 flow rate. ……………………………….………………………54 Figure 3-5b. H2O2 release profile from a 3.5 : 1 SR-19 form during a column leaching test using a 7 ml min-1 flow rate. ………………………………….……………………54 Figure 3-5c. H2O2 release profile from a 3.5 : 1 SR-20 form during a column leaching test using a 7 ml min-1 flow rate. ……………………………………………………….55 Figure 3-5d. H2O2 release profile from a 3.75 : 1 SR-21 form during a column leaching test using a 7 ml min-1 flow rate. ……………………………………………………….55 Figure 3-5e. H2O2 release profile from a 3.75 : 1 SR-22 form during a column leaching test using a 7 ml min-1 flow rate. …………………………………….…………………56 Figure 3-5f. H2O2 release profile from a 4 : 1 SR-23 form during a column leaching test using a 7 ml min-1 flow rate. ……………………………………………………………56 Figure 3-5g. H2O2 release profile from a 4 : 1 SR-24 form during a column leaching test using a 7 ml min-1 flow rate. …………………………………...……………………….57 Figure 3-5h. H2O2 release profile from a 4 : 1 SR-25 form during a column leaching test using a 7 ml min-1 flow rate. ………………………………………...………………….57 Figure 3-5i. H2O2 release profile from a 4 : 1 SR-26 form during a column leaching test using a 7 ml min-1 flow rate. ……………………………………...…………………….58 Figure 3-5j. H2O2 release profile from a 4 : 1 SR-27 form during a column leaching test using a 7 ml min-1 flow rate. ………………………………………..…………………..58 Figure 3-5k. H2O2 release profile from a 4 : 1 SR-28 form during a column leaching test using a 7 ml min-1 flow rate. …………………………………………...……………….59 Figure 3-6a. Average release profile of SR-HP 23-28 with one standard deviation error bars. ……………………………………………………………………...………………60 Figure 3-6b. Average release profile of SR-HP 23-28 with one standard deviation error bars. ……………………………………………………………..……………………….61 Figure 3-7. Peak H2O2 flux versus oxidant : resin mixing ratio for all SR-HP forms..…63 Figure 3-8. Peak H2O2 flux versus surface area (cm2) for all SR-HP forms………….. 63 Figure 3-9a. Ferrous iron removal efficiencies after 1 minute at different [Fe+2]/[H2O2] ratios…………………………………………………………………………….………..66 Figure 3-9b. Ferrous iron removal efficiencies after 5 minutes at different [Fe+2]/[H2O2] ratios………………………………………………………………………….…………..66

12 Figure 3-9c. Ferrous iron removal efficiencies after 10 minutes at different [Fe+2]/[H2O2] ratios ……………………………………………………………………67 Figure 3-10. Sketch map of BH006900 in relation to Brush Fork and sampling sites…..70 Figure 3-11. Dissolution of SR-HP forms varies depending on the heterogeneous nature of mixing. SR-HP forms with little resin on exterior exhibit much quicker dissolution..71 Figure 3-12a. Percent of Fe2+ removed during the field demonstration (normal time scale). ……………………………………………………………………………………72 Figure 3-12b. Percent of Fe2+ removed during the field demonstration (log time scale)……………………….. …………..…………..…………..………….……………73 Figure 3-13. Site BH00690 looking upstream, location of SR-HP treatment forms...…..74 Figure 3-14. Iron precipitating after one day of treatment at sample site #3 (24 meters downstream of mixing point. …………..…………..…………..………………………..74 Figure 3-15a. pH versus time measured at six different sampling locations (normal time scale). ………..…………..……………..…………..……………..………….....……….75 Figure 3-15b. pH versus time measured at six different sampling locations (log time scale). .…………..…………….. .…………..…………….. .…………..…...…………..75 Figure 3-16. pH versus distance from BH006900 measured at different times……...…..76 Figure 3-17a. Conductivity versus time measured at the six different sampling locations ……………………………………………………………………………………………77 Figure 3-17b. Conductivity versus time measured at the six different sampling locations. ……………………………………………………………………………………………77 Figure 3-18. Conductivity versus distance from BH006900 measured at different times. ……………………………………………………………………………………………78 Figure 3-19a. Fe2+ versus time measured at five sampling locations downstream of BH006900. …………………………………………………………………...………….79 Figure 3-19b. Fe2+ versus time measured at five sampling locations downstream of BH006900. …………………………………………………………………..…………. 79 Figure 3-20. Fe2+ versus distance from BH006900 measured at different times..….…..80 Figure 3-21. H2O2 versus time measured immediately upstream of BH006900…..……81 Figure 3-22a. Fe2+ measured at sampling location 2 and H2O2 fluxes versus time……81 Figure 3-22b. Fe2+ measured at sampling location 2 and H2O2 fluxes versus time...….82 Figure 3-23. [Fe2+ ]/[H2O2] ratio versus time. …………………………………………82

13 CHAPTER 1: INTRODUCTION

Sulfur-rich, acidic wastewaters are generated world-wide from a multitude of natural and anthropogenic sources. Acid mine drainage (AMD), although naturally occurring as part of the rock weathering process, is exacerbated by large disturbances made to the earth typically involved with metal and coal mining operations. Waters interacting with active and, notably, abandoned mines are often net acidic and pose environmental threats due to high concentrations of metals (e.g. iron, aluminum and manganese) (Johnson, 2005). Also referred to as acid rock drainage (ARD), AMD is primarily a product of the local mineralogy of rock material and the availability of water and oxygen to rapidly oxidize sulfide bearing minerals, notably pyrite (FeS2), the most abundant sulfide mineral on the planet. Characterized by low pH and high concentrations of metals and other toxic elements, AMD has potential to severely contaminate surface and groundwater (Akcil, 2006).

Coal extraction in the Appalachian region of the United States remains a stable industry today, building on a three-century legacy; however, adequate environmental regulations began appearing in policies less than 30 years ago. The Appalachian region, divided into three geographic provinces (Figure 1-1), has a long history of bituminous coal mining in particular the northern and central coal regions which account for 32% and

63% of the basin’s total production respectively (USGS, 2002). Unregulated coal mining in this region, although difficult to assess accurately, has left a significant environmental burden on the surrounding streams. The National Stream Survey (NSS), conducted by the U.S. Environmental Protection Agency, suggests that over 10,000 km of streams within the Appalachian region are impacted by AMD (Herlihy et al., 1990). 14 Several techniques exist for reducing AMD in impacted streams and lakes

(Johnson, 2005); however, the large number of contaminated sources and difficulty to characterize the sources poses a persistent challenge. Treatment options fall under two regimes: biological and abiotic, both having active and passive strategies.

Figure 1-1. Locations of the three Appalachian Basin coal regions (from USGS, 2002).

Active biological processes rely on the abilities of microorganisms to generate alkalinity, including denitrification, methanogenesis, sulfate reduction and iron and manganese reduction, subsequently immobilizing metals and reversing the reactions responsible for the production of AMD. Aerobic wetlands provide a passive biological method for oxidizing ferrous iron to ferric iron; however, this strategy is typically only constructed to treat net alkaline mine waters, because the main reaction is net acid- 15 generating (Johnson, 2005). Oxidation of iron at neutral pH values occurs rapidly; however, at pH values <4, the rate of iron oxidation is effectively independent of pH

(Stumm and Morgan, 1981).

The most frequent abatement strategy for mitigating acidic wastewaters is an active treatment process involving the addition of a neutralizing agent (Coulton et al.,

2003). Common neutralizing reagents used include calcium oxide (lime), slaked lime, calcium carbonate, sodium carbonate, sodium hydroxide, and magnesium oxide and hydroxide (Johnson, 2005). The added alkaline material raises the pH, accelerating the rate of oxidation of ferrous iron, subsequently causing metals in solution to precipitate as hydroxides and carbonates forming an iron-rich sludge. Neutralizing agents range in price and effectiveness depending on local water geochemistry. Although active treatment can effectively remediate AMD, disadvantages include high operating costs and problems disposing the iron-rich sludge. Alternatively, alkalinity can be added to

AMD through passive anoxic limestone drains (ALD). The goal of these systems is to avoid oxidation of ferrous iron and precipitation of ferric iron in the limestone, while maintaining alkali addition to the wastewater (Kleinmann et al., 1998).

Applicability of the different treatment systems is dependent on multiple variables, particularly cost, available space and local geochemistry of the mine water. For this study a new passive treatment system was designed using the slow release of oxidants from matrix-style diffusion. Recent studies have been conducted using slow release systems (SR) with advanced oxidation processes (AOPs) for treating organic pollutants in urban storm runoff (Sun, 2011; Tong, 2013; Eyerdom, 2014) and DNAPLs in groundwater (Lee and Schwartz, 2007a; Lee et al., 2008; Olson, 2011; Gupta, 2013). For 16 this study, several AOPs were evaluated for treating AMD water. Sodium percarbonate

(Na2CO3·1.5H2O2) was chosen as the reagent for developing a SR system because it readily dissolves in water (solubility = 150 g/L at 25°C) and release sodium carbonate and hydrogen peroxide (H2O2), a strong oxidant.

2(Na2CO3·1.5H2O2)  2 Na2CO3 + 3 H2O2 (l) (Eq. 1)

A SR approach to treat AMD impacted waters involves strategic placement of an oxidizing agent that is released at predetermined rates over time to serve as a passive treatment method over a long period of time. The treatment chemicals in the SR system create continual zones of oxidation and acid neutralization at a predetermined distance downstream of the installation point away from an underground or surface mine seep, reducing the sacrifice zone in which dissolved metals begin to precipitate. This passive treatment system is intended to work in a variety of conditions which allows for installation in remote areas without access to a power source. Slow-release forms could be installed for a predetermined amount of time from days to months based on treatment conditions and desired release rates and durations.

1.1 Formation of AMD

Waters impacted by AMD often yield high specific conductivity, high concentrations of iron, aluminum, and manganese, and have low pH. A significant amount of AMD is left untreated due to inadequate or expensive treatment options. The oxidation of pyrite (or other sulfide bearing minerals) best provides information on acid generation. The first step in this reaction involves the oxidation of the sulfide mineral, pyrite as an example, into dissolved ferrous iron, sulfate and hydrogen: 17 2+ 2− + FeS2 + 7/2 O2 + H2O  Fe + SO4 + 2 H (Eq. 2)

This initial reaction increases total dissolved solids and acidity of the water, leading to a reduction in pH unless neutralized by an alkaline source. Oxygen is a major rate-limiting factor in this step, deriving 87.5% of the sulfate’s oxygen from molecular oxygen and the remaining 12.5% from water (Taylor et al., 1984). Under oxidizing conditions, much of the ferrous iron will be further oxidized to ferric iron:

2+ + 3+ Fe + 1/4O2 + H → Fe + 1/2H2O (Eq. 3)

Ferric iron precipitates as Fe(OH)3 and jarosite between pH values 2.3 and 3.5, leaving little Fe3+ in solution and lowering pH as a result:

3+ + Fe + 3 H2O  Fe(OH)3 + 3 H (Eq. 4)

Ferric iron remaining in solution from Eq. (3) that does not precipitate in Eq. (4) may go on to further oxidize additional sulfide minerals, acting as the primary rate-limiting factor:

3+ 2+ 2− + FeS2 + 14Fe + 8H2O  15Fe + 2SO4 + 16H (Eq. 5)

In Eq. (5) 100% of the oxygen in the sulfate is derived from water. A combination of Eqs.

(2), (3) and (4) provides a simplified basic reaction of the oxidation of pyrite and the subsequent formation of acid and iron precipitate:

2- + FeS2 + 15/4O2 + 7/2 H2O  Fe(OH)3 + 2SO4 + 4 H (Eq. 6)

Several factors play a key role in the generation of acid from mine waste including pH, temperature, oxygen concentrations in both air and water, degree of saturation with water, chemical activity of Fe3+, bacterial activity, chemical activation energy required to initiate acid generation and surface area of exposed metal sulfides

(Akcil and Koldas, 2006).

18 1.2 Advanced Oxidation Reactions

Previous efforts have been made to create SR forms for treating organic pollutants in urban runoff (Tong, 2013; Eyerdom, 2014) utilizing AOPs. Specific oxidants used in these studies were persulfate reacting with ferrous iron (Fe2+) and hydrogen peroxide

(H2O2). These oxidants generate relatively quick reaction times and are colorless, odorless, and have high redox potentials. Oxidation by H2O2 alone is often not effective for high concentrations of organic contaminants because of low rates of reaction.

Transition metal salts can activate H2O2 to produce hydroxyl radicals which have higher redox potentials and yield Fenton’s reaction (Table 1-1). Under these conditions,

Fenton’s reagent is one of the most powerful oxidizers known. For this experiment

Fenton’s reaction was the primary AOP under investigation because Fe2+ is naturally available in AMD.

Table 1-1. Redox potential of common oxidants Oxidant Redox potential E0 (V)

Hydroxyl radical OH● 2.76

ozone O3 2.07

Hydrogen peroxide H2O2 1.78

● Perhydroxyl radical HO2 1.7

19 1.2.1 Fenton’s Reagent

Hydrogen peroxide activated by Fe2+ has proven to be an effective oxidizer for treating coal mining residue (Silva, 2011) and organic compounds (Chamarro, 2000).

Hydrogen peroxide can be added into solution through the dissolution of sodium percarbonate (Na2CO3·1.5H2O2), a commonly used commercial and industrial detergent

(Eq. 7). The ability to add a strong oxidizer and carbonate species for acid neutralization makes sodium percarbonate a good candidate for AMD remediation.

2(Na2CO3·1.5H2O2)  2 Na2CO3 + 3 H2O2 (l) (Eq. 7)

2+ Fenton’s reaction utilizes an iron (Fe or zero-valent iron) catalyst to convert H2O2 to free radicals with high redox potentials. Fenton’s reaction catalyzed by Fe2+ are described as Eqs. 8 and 9 (Neyens and Baeyens, 2003):

2+ 3+ ● - Fe + H2O2  Fe + OH + OH (Eq. 8)

k= 70 M-1s-1 (Rigg et al.,1954)

OH● + Fe2+  OH- + Fe3+ (Eq. 9)

k= 3.2 * 108 M-1s-1 (Buxton and Greenstock, 1988)

3+ Ferric iron (Fe ) produced in Eqs. 8 and 9 can further catalyze H2O2, forming water and oxygen. Additional ferrous ions and radicals are formed in Eqs. 10-13 (Neyens and

Baeyens, 2003).

3+ 2+ + Fe + H2O2  Fe-OOH + H (Eq. 10)

k= 0.001-0.01 M-1s-1 (Walling and Goosen, 1973)

2+ ● 3+ - Fe + HO2  Fe + HO2 (Eq. 11)

k= 1.3 * 106 M-1s-1 (at pH=3, Bielski et al., 1985)

3+ ● 2+ + Fe + HO2  Fe + O2 + H (Eq. 12) 20 k=1.2 * 106 M-1s-1 (at pH=3, Bielski et al., 1985)

● ● OH + H2O2  H2O + HO2 (Eq. 13)

k= 3.3 * 107 M-1s-1 (Buxton and Greenstock, 1988)

A simplified version of Fenton’s reaction which takes into account the dissociation of water can be observed in Eq. 14 (Walling, 1975).

2+ + 3+ 2Fe + H2O2 +2H  2Fe +2H2O (Eq. 14)

+ The reaction in Eq. 14 implies that H is required to decompose H2O2, indicating the need for acidic environments for maximum Fenton reaction efficiency. Previous studies have suggested pH values close to 3 for optimum hydroxyl radical formation (Neyens and

Baeyens, 2003).

1.3 Slow Release Systems

Slow release (SR) systems were initially developed as a way to maintain drug levels within a certain window to avoid high concentrations and maximize therapeutic efficiency (Hoffman, 2008). These systems allow chemicals to be released at a predetermined rate for a definite amount of time (days to years) (Lee and Schwartz,

2007a). Similar technology has been used in agriculture to control the release of nutrients from fertilizers (Du et al., 2006). Recently, utilization of SR systems have expanded to environmental remediation. In moving water, SR forms can release oxidizing chemicals at predetermined rates that react with contaminants, breaking them down to non-harmful products.

Two main types of SR systems are commonly used: encapsulation and matrix styles. Encapsulation SR systems are manufactured with a treatment chemical inside of a 21 single-layered polymeric coating. These styles tend to yield shorter release durations, as only one-layered coating is available and the chemical is dumped when the coating breaks. Matrix SR systems are a better choice for environmental remediation, yielding longer and more stable and controlled release rates. In simplest matrix system, i.e., monolithic SR systems, treatment chemicals are mixed homogeneously within a matrix.

As the matrix dissolves, secondary porosity is created, which allows the chemicals inside the matrix to slowly dissolve based on diffusion (Figure 1-2).

Figure 1-2. Slow release diffusion modeled in cross sectional view. (Lee and Schwartz, 2007a).

Lee et al. (2008) demonstrated the effectiveness of matrix SR using KMnO4 in the destruction of TCE with column leaching experiments, proof-of-concept flow-tank tests, and model simulations. The researchers were able to create a long-term stable zone of oxidation with somewhat constant permanganate levels. 22 For this study, knowledge of matrix SR systems were applied to treat acid mine drainage using H2O2. Ferrous iron commonly available with acidic mine water acts as a catalyst to produce hydroxyl radicals that are powerful oxidizers.

1.4 Study Site

Four potential field sites were established within the Monday Creek watershed

(303 km2), northwest of Athens, OH. During the late 19th century, settlers entered this region and began to mine extensive amounts of coal both on the surface and underground.

Underground mines pierced to the Middle Kittanning (No. 6) Coal of the Allegheny

Series, Pennsylvanian System and mined with a room-and-pillar system averaging 2.0 to

2.3 meters thick (Pigati and Lopez, 1999). Data provided by the Monday Creek

Restoration Project is available through Ohio University’s Voinovich School of

Leadership and Public Affairs’ non-point source (NPS) monitoring project.

After initial background chemical sampling, site BH00690 was selected as the pilot study site due to high concentrations of dissolved iron and low perennial flow (<2

L/sec). Acid mine drainage produced at this site discharges into Brush Fork before draining into Snow Fork, one of the main tributaries of Monday Creek. This site is located approximately 5 miles due north of Nelsonville, Ohio within the Wayne National

Forest (Figure 1-3). 23

Figure 1-3. Location of study site BH00690 within the Monday Creek Watershed (adapted from Monday Creek Reclamation Project).

Figure 1-4. Location of study site BH00690 in relation to Jobs New Pittsburg Rd and Brush Fork, and other AMD sampling points stored on NPS website (http://watersheddata.com)

With each of these sites a clean stream, not significantly impaired by AMD, converges with a seep discharging acidic mine water. Previous data stored on the NPS website provided background information on the site (Table 1-2); however, incomplete data sets will require more comprehensive sampling technique.

This study focused on forming Fenton’s reagent to oxidize and remove dissolved ferrous iron from the system. As a result of sodium percarbonate 2(Na2CO3·1.5H2O2) dissolution, alkinity will be added and pH will be affected. If the pH increases to a certain level, metals will begin to precipitate depending on their chemical properties

(Figure 1-5). Wei et al. (2005) found that 75% of aluminum precipitated as the pH increased from 3.5 to 4.5 and almost all was removed 5.0 < pH < 8.0 while manganese requires a higher pH value for precipitation (>8.5). 25

Figure 1-5. Solubility of major and minor metals at different pHs (modified from Wei et al., 2005)

26 Table 1-2. Historic water chemistry and discharge measurements at study site BH00690. Acidi Alkali D.O. TDS Discharg Conductivity ty nity Date Temp (°C) (mg/ pH (mg/ e (L/sec) (μS/cm) (mg/ (mg/L L) L) L) ) 10/7/2002 12 1.27 3.03 2070 553 0 1820 5/24/2010 21.9 1.16 3.6 3.07 1870 444 0 1520 7/21/2010 18.4 1.90 6 3.14 1910 465 0 1590 8/16/2010 11.7 2.10 2.7 3.98 1780 473 0 1690 9/21/2010 12.3 0.93 2.2 3.66 1770 468 0 1740 11/1/2010 9.3 0.79 3.3 3.66 1820 490 0 1740 11/22/2010 7.9 0.54 2 3.38 1890 486 0 1680

Ca Mg Mn Al Total Fe Dissolved Ferrous Sulfate Date (mg/ (mg (mg/ (mg/ (mg/L) Fe (mg/L) Fe (mg/L) (mg/L) L) /L) L) L) 10/7/2002 203 207 182 1194 4.83 14.2 5/24/2010 187 186 173 1007 122 55.9 3.58 4.16 7/21/2010 201 196 120 962 121 55.7 3.42 3.98 8/16/2010 247 252 244 1005 129 58.2 3.51 4.55 9/21/2010 249 240 232 1037 131 59.2 3.55 4.41 11/1/2010 248 244 228 1033 132 59.4 3.64 4.64 11/22/2010 246 252 207 1138 138 63 3.72 5.2

1.5 Study Objectives

This study will focus on developing CRS forms that can release hydrogen peroxide (H2O2) and carbonates (SR-HP) at a predetermined rate over an extended amount of time. The SR-HP forms can be installed in streams not impacted by AMD, but combines with AMD affected streams. Ferrous iron in the AMD stream will catalyze

2+ H2O2 to create Fenton’s reaction, which can oxidize Fe in the AMD streams. Specific objectives include (i) preliminary hydro-chemical investigations of potential future field demonstration sites, (ii) development and characterization of SR-HP forms through column tests, (iii) proof-of-concept tests for estimating the efficacy of the SR-HP form in 27 oxidizing metals and neutralizing acids in AMD, and (iv) small-scale proof-of-concept field investigations using the designed SR-HP forms.

28 CHAPTER 2: MATERIALS AND METHODS

2.1 Materials

Deionized water was produced with a Milli-Q system from Millipore. Sodium percarbonate (Na2CO3·1.5H2O2), a colorless, free-flowing, granular source of anhydrous

H2O2, was purchased from The Chemistry Store and used for manufacturing SR forms.

This adduct of sodium carbonate and hydrogen peroxide readily dissolves in water

(solubility = 150 g/L at 25°C; Jones, 1999) (Eq. 15). Soda ash produced in the dissolution of Na2CO3·1.5H2O2 is widely used as a relatively strong base in many AMD treatment projects to increase alkalinity. The ability for Na2CO3·1.5H2O2 to produce both a strong oxidant and base makes it a valuable chemical treatment option for both oxidation and neutralization of AMD. Polyester casting resin, used as a binding agent in

SR forms, was purchased from Dick Blick art supplies.

2(Na2CO3·1.5H2O2)  2Na2CO3 + 3 H2O2 (Eq. 15)

Ethanol (5% 2-Propanol and 5% methanol, Acros Organics), Neocuproine (99+%,

Acros Organics), and cupric sulfate (1%) were purchased from Fisher Scientific and used for hydrogen peroxide concentration determination. Hydroxlamine (NH2OH*HCl, 99+%, from Acros Organics), 1,10-phenanthroline monohydrate (99+%, Acros Organics), and

HCl (1N, Acros Organics) were purchased from Fisher Scientific for Fe2+ and total Fe concentration determination. Ferrous iron was available from locally collected AMD.

2.2 Field Measurements

At site BH00690, field data was collected and compared utilizing literature available from the NPS Monitoring Project created by the Voinovich School of 29 Leadership and Public Affairs at Ohio University. Data sets included field chemistry, grab samples, flow measurements, and distance from chemical addition following water sampling protocols established by Ohio EPA (2003). Field chemistry, consisting of dissolved oxygen (DO), oxidation-reduction potential (ORP), pH, conductivity, temperature, were measured using a YSI 600 XLM data sonde (ODNR).

Grab samples were collected with a stainless steel bucket at mid-depth after three rinses of water from the sampling site. Sampled water was analyzed at Ohio Department of Natural Resources (ODNR) for complete Group II parameters including, Na, K, Si, and Cl concentrations following the appropriate standard method for each parameter using American Society for Testing and Materials (ASTM) 2005 guidelines.

Flow measurements were made with a fixed or portable flume when discharge was between 2.5 – 1000 gallons / minute. A bucket-and-stopwatch method was utilized to measure flows less than 10 gallons / minute. All field and sampling devices were provided by ODNR.

2.3 Designing Slow Release Forms

Previous implementation of the SR system involved embedding granular reactive oxidants in polymeric matrices made from paraffin wax (Tong, 2013; Eyerdom, 2014).

Dissolution of the matrix becomes diffusion controlled due to porosity developed in the polymer. Chemical release can then be controlled by choice of oxidants with different solubilities, density of the treatment chemical in matrix, amount of mixing, granule size, and the overall size of the SR forms. 30 For designing a SR system for AMD, experiments that optimize oxidant-release and alkali-release agents provided desired release rates. Column leaching experiments were used to measure the release rate for various potential SR forms following experiments conducted by Lee and Schwartz (2007a). Slow release forms were made using reagent grade sodium percarbonate (Na2CO3·1.5H2O2) and casting resin as a matrix.

Casting resin was chosen over paraffin wax to create more durable SR forms that could contain a higher percentage of oxidant compared to SR forms made with paraffin wax.

Sodium percarbonate and casting resin were measured separately to insure proper mixing ratios and then homogeneously mixed together using a blender. A few drops of catalyst must be added to the polyester casting resin in order for the forms to harden completely.

To increase the release rates of SR forms, holes were drilled in six of the forms ranging from 1/16” to 1/8” in order to increase the surface area in which diffusion occurred.

Surface area was calculated by measuring the diameter and height of the original SR-HP form and the holes drilled into the form. Surface area created from the cylindrical holes was added to the original calculated surface area.

2.4 Estimating Release Rate: Column Leaching Tests

Slow release forms were submerged with deionized water in glass columns (L x

D : 15 cm x 4.8 cm) to measure the release rates. Peristaltic pumps maintained a determined flow rate (7 to 10 mL min-1) through the column to obtain a perfect sink condition. Release rates were calculated from different oxidant : matrix ratios. Effluent samples (8 mL min-1) were collected at fixed time intervals. Further spectrophotometric analysis was used to determine mass fluxes and release rates. Preliminary results from 31 Tong (2013) on the use of SR-HP for treating organic pollutants in urban runoff suggests that wax based SR-HP can release H2O2 over an extended period of time in a somewhat controlled manner. Control release forms with higher oxidant : wax mixing ratios yielded higher release rates (Figure 2-1). However, the reproducibility of the SR-HP forms in terms of their release kinetics and durations was not evaluated. Furthermore, previous results from Tong (2013) and Eyerdom (2014) produced low recovery rates of hydrogen peroxide based on mass-balance stoichiometry and release rates. This study focused on developing SR forms that yield consistent release kinetics and durations at variable mixing ratios and attain greater recovery rates of hydrogen peroxide. In addition, casting resin was used as a matrix material, instead of paraffin wax, to generate more durable SR forms with stable and higher release rates.

Figure 2-1: Release kinetics of SR-HP forms at 8 mL min -1 column test (adapted from Tong, 2013) 32

Sustained release rates from SR-HP column tests were determined by evaluating the release profiles. The H2O2 fluxes from column tests reach a peak release rate within the first several hours of the experiments. The slope of the release profile becomes much smaller after 4 hours. These stabilized release values were modeled using a first order decay equation where A is a constant and k represents the decay coefficient (Eq. 16).

Computing t-statistics from A and the standard error associated with the model provided a way to compare the consistency of the stabilized release among different forms.

Flux = Ae-kt (16)

2.4.1 Determination of Hydrogen Peroxide Concentration

Hydrogen peroxide concentrations were determined following the spectrophotometric method from Baga et al. 1988. This procedure is based on the reduction of copper (II) ions by hydrogen peroxide in the presence of excess 2, 0- dimethyl-1 , 10- phenanthroline (DMP) to form the copper (I)- DMP complex (Eq. 17).

The copper (I) – DMP complex was determined directly by spectrophotometric measurement at 454 nm (Baga et al., 1988).

2+ 2+ + 2Cu + 2DMP + H2O2  2Cu (DMP) + O2 + 2H (17)

DMP solution is formed by dissolving 1 g of neocuprione into 100 mL of ethanol.

A 0.01 mol/L Cu (II) sulfate stock solution is made by dissolving 25 mL 10 g/L cupric sulfate into 75 mL of deionized water. A mixture of 1 mL Cu (II) sulfate stock solution,

1 mL DMP solution, and 8 mL of H2O2 sample solution are measured for absorbance at 33 454 nm in UV-VIS spectrophotometer. Blank samples were used often to calibrate the

UV-VIS. The calibration curve used for this method is shown in Figure 2-2.

1.8

1.6

1.4

1.2

0.8 Absorbance 0.6

0.4

0.2

0 0 1 2 3 4 5 6 Concentration (mgL-1)

Figure 2-2. Standard curve for Hydrogen Peroxide concentration determination

2.4.2 Determination of Ferrous Iron Concentration

Ferrous iron concentrations were determined using a UV-VIS spectrophotometer based on the reaction between ferrous iron and 1, 10-phenanthroline at pH 2~9 (Eq. 18).

+2 The production complex [Fe(C12H8N2)3] can be determined by spectrophotometric measurement at 510 nm. Hydroxylamine was used to reduce Fe3+ before total Fe concentration determination (Eq. 19). These methods are used to determine Fe2+ and total Fe concentrations in aqueous solutions between 0.1 to 6.0 mg L-1. Samples were diluted to meet this concentration range requirement for analysis. 34 2+ +2 Fe + 3C12H8N2  [Fe(C12H8N2)3] (Orange) (Eq. 18)

4FeCl3 + 2NH2OH·HCl  4FeCl2 + N2O + 6HCl + H2O (Eq. 19)

One mL of acetate buffer is mixed with 1 mL of water sample in a 40 mL EPA glass vial. For Fe2+ concentration determination, one mL 1, 10-phenanthroline is added to the vial. For total Fe concentration determination, 0.5 mL 1, 10-phenanthroline and

0.5 mL 10% hydroxylamine solution is added to the glass vial. Each solution is hand shaken and requires a reaction time of 15 minutes to equilibrate. Concentrations of Fe2+ were calibrated from Fe standard solution (Figure 2-3).

3.5

2.5

1.5 Absorbance

0.5

0 0 10 20 30 40 50 60 Concentration (mgL-1)

Figure 2-3. Standard curve for Ferrous Iron concentration determination

35 2.5 Proof-of-Concept Ferrous Iron Removal Test

Experiments were performed in the laboratory to simulate Fe2+ removal with SR-

HP forms using Fenton’s reagent. Nineteen tests were conducted using different

2+ concentrations of both Fe and H2O2 to measure the removal efficiency for different

+2 [Fe ]/[H2O2] ratios. Following research results from Mahiroglu and others (2009), which used the Fenton process to treat acid mine drainage from a copper mine, initial

+2 [Fe ]/[H2O2] ratios approached 1.8 for optimal removal efficiency. Final ratios of

+2 [Fe ]/[H2O2] ranged from 0.5 to 10.5.

Hydrogen peroxide concentrations from SR-HP column leaching tests were measured using a spectrophotometer following the DMP method (Baga et al., 1988).

Total and Fe2+ concentrations, from acid mine drainage collected at one of the potential study sites, were determined with a spectrophotometer following previously described methods. After determining solution concentrations, volumes were measured to assure

+2 predetermined [Fe ]/[H2O2] molar ratios would be attained. The reactions occurred

2+ rapidly, lasting up to fifteen minutes long. During the reaction, H2O2, total and Fe concentrations were measured in addition to pH and conductivity.

2.6 Small-scale Field Application Test

The NPS Monitoring Project, operated by Ohio University’s Voinovich School of

Public Affairs, contains data collected by watershed groups and the Ohio DNR from the mid-1990s. Review of these data was used for quantifying the ranges in metal concentrations to manufacture the most effective SR-HP forms. Slow release forms that release the minimum required oxidants for oxidizing metals in the AMD-affected stream 36 were prepared. For simplicity, this study focused on treating iron. Iron concentrations and discharge data were used for determining the minimum release rates. The SR-HP forms were placed in an unimpacted stream that flows into an AMD seep (BH00690).

Metal concentrations and pH values measured upstream and downstream of the reaction zone before and after treatment were used to evaluate metal removal and acid neutralization efficiency.

Six sites were established at the field study area to collect samples and analyze for ferrous iron concentrations. Site 1 captured the release of SR-HP, immediately downstream of the installed forms, but upstream of any AMD to ensure that the forms would not become armored with iron. Background iron concentrations were negligible

-1 here (<1 mg L ) and only H2O2 concentrations were analyzed at this site. Sites 2-6 were spaced out evenly in 25 foot intervals downstream of the confluence between the AMD and clean water. Hydrochloric acid was added to these samples to prevent ferrous iron from oxidizing before laboratory analysis. Sodium thiosulfate was also added at sites 2-6

2+ to quench the reaction between H2O2 and Fe , stopping the reaction (Keen et al., 2013).

37 CHAPTER 3: RESULTS AND DISSCUSION

3.1 Study Site

Initial background sampling was performed on June 10th, 2014 with the aid from the Monday Creek Restoration Project. Four sites within the Monday Creek watershed were targeted for a small-scale field demonstration and sampled for field chemistry, grab samples, and flow measurements. Site BH00690 was selected as the pilot study site due to low perennial flow yet high concentrations of Fe2+ available for treatment and Fenton’s process to proceed (Figure 3-1). This upwelling mine pool had been sampled 9 times, from 2001 – 2010, by the Monday Creek Restoration Project and Ohio Department of

Natural Resources (Table 3-1). Historical and measured stream discharge and Fe2+ concentrations were used to determine the number of SR forms to install at the site for efficient iron removal.

Figure 3-1: Small-scale field demonstration site, BH006900, located within the Monday Creek watershed. Baseline chemical sampling on June 10th, 2014 38 Table 3-1. Historic discharge and chemical parameter measurements at BH00690 (http://watersheddata.com) Discharge Conductivity Acidity Fe (tot) Fe 2+ Date (L s-1) ORP pH (µS cm-1) (mg L-1) (mg L-1) (mg L-1)

5/24/2010 1.16 379.9 3.07 1870 444 187 173

7/21/2010 1.90 404.8 3.14 1910 465 201 120

8/16/2010 2.10 213 3.98 1780 473 247 244

9/21/2010 0.93 -6 3.66 1770 468 249 232

11/1/2010 0.79 90 3.66 1820 490 248 228

11/22/2010 0.54 393.9 3.38 1890 486 246 207

6/10/2014 0.61 466 2.82 1810 327 94 56.6

Average 2.81 277.37 3.37 1845.00 461.40 217.70 188.07

Median 1.22 379.90 3.42 1830.00 468.00 228.00 207.00

3.1.1 Baseline Chemical Sampling

Field chemistry consisted of temperature, oxidation-reduction potential (ORP), pH, specific conductivity, and flow rate (Table 3-2). Grab samples collected contained

Group II ions including Na, K, Cl, Fe2+ and total iron concentrations (Table 3-2). Results from Table 3-1 indicate the flow recorded on June 10th, 2014 (0.606 Lsec-1) was about half of the historic median (1.22 Lsec-1). Total and Fe2+ concentrations were also much lower than the historic medians, 94.0 and 56.6 mg L-1 compared to 228.0 and 188.07 mg

L-1 respectively. This site was chosen as the small-scale field demonstration location due to the high concentration of Fe2+ and low discharge, which would require less SR-HP forms to oxidize and remove the Fe2+. Historic dissolved oxygen at this site was low (3.3 mg L-1), suggesting most of the iron present at this site is in the more reduced form, Fe2+.

Background chemical sampling was useful, in combination with SR-HP column test 39 results, to calculate how many SR-HP forms would be required for efficient iron treatment at the site.

Table 3-2. Aqueous chemistry data for BH006900 sampled on 6/10/2014

Parameter Values Temperature (°C) 16.0 Discharge (L s-1) 0.606

Dissolved Oxygen (mg L-1) n.a. ORP 466 pH 2.82 -1 Conductivity (µS cm ) 1810 Acidity (mg L-1) 327 Alkalinity (mg L-1) 0.00 Sulfate (mg L-1) 691 Cl (mg L-1) 6.53 Ca (mg L-1) 94.2

Mg (mg L-1) 40.7 -1 Na (mg L ) 12.2 K (mg L-1) 7.54 Mn (mg L-1) 2.61 Al (mg L-1) 1.33 Total Fe (mg L-1) 94.0 Dissolved Fe (mg L-1) 91.8 -1 Ferrous Fe (mg L ) 56.6 -1 Total Fe Load (kg day ) 4.92 Ferrous Fe Load (kg day-1) 2.96 -1 TDS (mg L ) 1060 TSS (mg L-1) 7.00 Hardness (mg L-1) 402

40 3.2 Characterizing Slow Release Systems

3.2.1 Slow Release Form Dimensions

Slow-release forms were designed using different oxidant : resin mixing ratios.

Mass, volume, and density were determined before and after the column leaching tests were performed. These data were then used to calculate the amount of available H2O2 before and after the column leaching experiments were performed, based on stoichiometric relationships (Table 3-3).

Table 3-3. Properties of slow-release forms made with sodium percarbonate and casting resin, shaded cells separates the different groups of oxidant : resin mixing ratios discussed in depth SR Mass Width Height Volume Density Mass Mass Recovery 3 -3 Form Ratio (g) (cm) (cm) (cm ) (gcm ) Oxidant (g) H2O2 (g) H2O2 (%) 1 1.7:1 73.5 2.8 9.0 55.4 1.33 45.9 14.9 0.47 2 2.5:1 64.3 2.8 7.8 48.0 1.33 45.9 14.9 1.6 3 3.3:1 58.9 2.8 7.5 46.2 1.28 45.3 14.7 48 4 5:1 55.5 2.7 7.2 42.8 1.3 46.2 15.0 35 5 5:1 55.5 2.7 7.2 42.8 1.3 46.2 15.0 26 6 6:1 61.7 2.8 8.2 50.5 1.22 52.9 17.2 21 7 6:1 66.0 2.8 9.5 58.5 1.13 56.6 18.4 14 8 5:1 66.9 2.8 8.5 52.3 1.28 55.8 18.1 34 9 5:1 77.8 2.8 11.5 70.8 1.1 64.8 21.1 53 10 5:1 81.9 2.8 12.0 73.9 1.11 68.2 22.2 54 11 4.2:1 75.6 2.7 8.8 52.6 1.44 61.0 19.8 19 12 3.8:1 75.1 2.7 8.6 51.1 1.47 59.5 19.3 9.1 13 3.5:1 87.2 2.7 10.3 61.2 1.43 67.0 21.8 47 14 3.7:1 82.1 2.7 10.0 59.4 1.38 64.6 21.0 57 15 3.5:1 83.1 2.7 9.9 56.7 1.47 64.6 21.0 50 16 3.5:1 87.2 2.7 10.3 59.0 1.48 67.8 22.0 42 17 3.75:1 94.8 2.7 11.0 63.0 1.5 74.8 24.3 52 18 3.75:1 93.8 2.7 11.4 65.3 1.44 74.1 24.1 62 19 3.5:1 69.1 2.8 10.0 46.2 1.5 53.8 17.5 50 20 3.5:1 72.3 2.8 9.3 48.4 1.49 56.2 18.3 44 21 3.75:1 71.5 2.8 12.3 47.1 1.52 56.4 18.3 76 22 3.75:1 75.4 2.8 13.1 50.6 1.49 59.5 19.3 63 23 4:1 87.4 2.8 12.9 63.6 1.37 70.0 22.7 68 24 4:1 91.0 2.8 13.8 69.9 1.3 72.8 23.7 65 25 4:1 98.6 2.8 14.3 72.6 1.36 78.9 25.6 64 26 4:1 96.8 2.8 13.4 75.5 1.28 77.4 25.2 81 27 4:1 96.6 2.8 14.8 76.9 1.26 77.3 25.1 50 28 4:1 102.7 2.8 14.0 72.5 1.42 82.1 26.7 45 41

Oxidant : resin mixing ratios ranged from 1.7:1 to 6:1, resulting in SR-HP forms comprising a range of H2O2 concentration. Sodium percarbonate is comprised of 34%

H2O2 per molecular weight, therefore, SR-HP forms ranged from 21 to 28% H2O2 depending on mixing ratio (Table 3-3). Recovery rate refers to the end measured mass of

H2O2 as a percentage compared to the calculated mass before the leaching tests began.

As the column leaching tests proceeded and Na2CO3·1.5H2O2 began to dissolve

(solubility = 140 g L-1), density would decrease. With greater oxidant : resin mixing ratio

(> 3:1), within a few days the SR-HP forms would begin to float in the columns. To prevent this from happening weights were added in the column with a plastic bottle.

3.2.2 Column Release Tests

Hydrogen peroxide concentrations were measured using a spectrophotometer during the column leaching tests. Combined with flow rate data, the flux of H2O2 was monitored during the tests in addition to total recovered H2O2 (Figures 3-2 – 3-5).

Peristaltic pumps, with a flow rate of 7-8 mL min-1, were used for all of the tests. The

-1 column tests were performed until H2O2 concentrations became negligible (< 1 mg L ).

Test durations lasted from days to months. Hydrogen peroxide concentrations were measured until samples were below the detection limit. Slow-release forms with lower oxidant : resin mixing ratios (< 3.5:1) trapped non-dissolved Na2CO3·1.5H2O2 within the matrix with too much casting resin, resulting in low H2O2 flux measurements and inefficient recover rates (Figures 3-2a-c). Slow-release forms with higher oxidant : resin mixing ratios (> 4:1) became structurally unstable and began to deteriorate inside the 42 columns within a few days. These tests resulted in high peak release of H2O2, but did not produce sustained lower releases (Figures 3-3a-g).

Figure 3-2a. H2O2 release profile from a 1.66 : 1 SR-1 form during a column leaching test using a 7 ml min-1 flow rate. 43

Figure 3-2b. H2O2 release profile from a 2.5 : 1 SR-2 form during a column leaching test using a 7 ml min-1 flow rate.

Figure 3-2c. H2O2 release profile from a 3.33 : 1 SR-3 form during a column leaching test using a 7 ml min-1 flow rate.

The first SR-HP forms were produced with little experience and resulted in low recovery rates (16.6% on average). Figures 3.2a-b (SR-1 and SR-2) produced the lowest peak release rates (0.05 and 0.112 mg min-1), sustained release rates (0.02 mg min-1) and recovery of H2O2 during the duration of the test compared to all other SR-HP column leaching tests. These two SR forms were manufactured with the lowest oxidant : resin mixing ratios (1.7 and 2.5 respectively), leading to much Na2CO3·1.5H2O2 being trapped by the polyester casting resin, inhibiting the dissolution of Na2CO3·1.5H2O2. Slow- release form 3 had a comparatively more efficient peak release rate (2.49 mg min-1),

-1 sustained release rate (0.2 mg min ), and H2O2 recovery rate (48%).

Figure 3-3a. H2O2 release profile from a 5 : 1 SR-4 form during a column leaching test using a 7 ml min-1 flow rate.

Figure 3-3b. H2O2 release profile from a 5 : 1 SR-5 form during a column leaching test using a 7 ml min-1 flow rate.

Figure 3-3c. H2O2 release profile from a 6 : 1 SR-6 form during a column leaching test using a 7 ml min-1 flow rate. 46

Figure 3-3d. H2O2 release profile from a 6 : 1 SR-7 form during a column leaching test using a 7 ml min-1 flow rate.

Figure 3-3e. H2O2 release profile from a 5 : 1 SR-8 form during a column leaching test using a 7 ml min-1 flow rate. 47

Figure 3-3f. H2O2 release profile from a 5 : 1 SR-9 form during a column leaching test using a 7 ml min-1 flow rate.

Figure 3-3g. H2O2 release profile from a 5 : 1 SR-10 form during a column leaching test using a 7 ml min-1 flow rate.

48 Slow release forms were produced with higher oxidant : resin mixing ratios to improve release rates and over efficiencies. Figures 3-3a-g show the release profiles of

SR-4-10, made with oxidant : resin mixing ratios < 4.5, the largest ratios manufactured in this experiment. Although these forms did have higher peak release rates (12.2 mg min-1),

-1 sustained release rates (0.36 mg min ), and recovery of H2O2 (34%) compared to SR- 1-3, these SR forms did not maintain structural integrity, thus would not be a good choice for field treatment with greater flows. These forms produced detectable H2O2 for less than 5 days on average.

Each test shows an initial peak in the release of H2O2 followed by a rapid decrease, producing stabilizing release rates from days to weeks, which is consistent with results from previous studies (Lee et al., 2008; Tong, 2013; Eyerdom, 2014). Normally, the magnitude of the initial peak is governed by the oxidant : resin mixing ratio. Higher ratios yielded higher peak release rates. Peak H2O2 release rates ranged between 5.0 x

10-2 mg min-1 and 52.1 mg min-1 with an average of 17.6 mg min-1. Stabilized release rates ranged between 2.0 x 10-2 and 1.5 mg min-1 with an average of 4.3 x 10-1 mg min-1.

Maximum available H2O2 for each SR-HP form was calculated based on stoichiometry and mass balance equations (Table 3-3). This amount was then compared to the total amount of recovered H2O2 based on the column leaching tests. Recovery rates of H2O2 for the SR forms varied from 0.47% to 81.4% due to a number of reasons.

One of major goals for this study was to obtain high recovery rate of H2O2 in SR-HP forms. By using polyester casting resin as a binding agent, higher oxidant : resin ratios were able to be used yet still produce structurally sound SR-HP forms, compared to 49 parrifin wax which required lower oxidant : wax mixing ratios to remain structurally intact.

The first SR-HP forms that were manufactured had lower oxidant : resin mixing

-1 ratios (< 3.5 :1) and yield the lowest peak H2O2 release rates (0.88 mg min ) and total recovery rates (16.6%)(Figure 3-2a-c). Hydrogen peroxide recovery efficiencies of SR-

HP forms began to increase as forms were manufactured with greater oxidant : resin mixing ratios, yet low enough to retain structural integrity (3.5 < oxidant : resin < 4.5)

(Figures 3-4 and 3-5). Slow-release forms made within this ratio range had higher initial release rates (22.5 mg min-1), stable release rates (5.2 x 10-1 mg min-1), recovery rates

(52.4%), and were able to produce sustained stable release rates for 14 days, on average.

Compared to SR-HP forms not manufactured within the ideal oxidant : resin range which, on average, had lower peak releases (8.8 mg min-1), stable release rates (2.7 x 10-1 mg min-1), recover rates (28.7%), and were able to produce sustained stable release rates for only about 7 days.

Figure 3-4a. H2O2 release profile from a 4.2 : 1 SR-11 form during a column leaching test using a 7 ml min-1 flow rate.

Some SR-HP column leaching tests did not experience a constant, stabilized decay. Hydrogen peroxide concentrations were seen to have multiple peaks as a result of pore dissolution allowing for secondary diffusion to increase H2O2 concentrations for a brief period (Figure 3-4b). 51

Figure 3-4b. H2O2 release profile from a 3.8 : 1 SR-12 form during a column leaching test using a 7 ml min-1 flow rate.

Figure 3-4c. H2O2 release profile from a 3.5 : 1 SR-13 form during a column leaching test using a 7 ml min-1 flow rate.

Figure 3-4d. H2O2 release profile from a 3.7 : 1 SR-14 form during a column leaching test using a 7 ml min-1 flow rate.

Figure 3-4e. H2O2 release profile from a 3.5 : 1 SR-15 form during a column leaching test using a 7 ml min-1 flow rate. 53

Figure 3-4f. H2O2 release profile from a 3.5 : 1 SR-16 form during a column leaching test using a 7 ml min-1 flow rate.

Figure 3-4g. H2O2 release profile from a 3.75 : 1 SR-17 form during a column leaching test using a 7 ml min-1 flow rate. 54

Figure 3-5a. H2O2 release profile from a 3.75 : 1 SR-18 form during a column leaching test using a 7 ml min-1 flow rate.

Figure 3-5b. H2O2 release profile from a 3.5 : 1 SR-19 form during a column leaching test using a 7 ml min-1 flow rate. 55

Figure 3-5c. H2O2 release profile from a 3.5 : 1 SR-20 form during a column leaching test using a 7 ml min-1 flow rate.

Figure 3-5d. H2O2 release profile from a 3.75 : 1 SR-21 form during a column leaching test using a 7 ml min-1 flow rate.

Figure 3-5e. H2O2 release profile from a 3.75 : 1 SR-22 form during a column leaching test using a 7 ml min-1 flow rate.

Figure 3-5f. H2O2 release profile from a 4 : 1 SR-23 form during a column leaching test using a 7 ml min-1 flow rate.

Figure 3-5g. H2O2 release profile from a 4 : 1 SR-24 form during a column leaching test using a 7 ml min-1 flow rate.

Figure 3-5h. H2O2 release profile from a 4 : 1 SR-25 form during a column leaching test using a 7 ml min-1 flow rate.

Figure 3-5i. H2O2 release profile from a 4 : 1 SR-26 form during a column leaching test using a 7 ml min-1 flow rate.

Figure 3-5j. H2O2 release profile from a 4 : 1 SR-27 form during a column leaching test using a 7 ml min-1 flow rate. 59

Figure 3-5k. H2O2 release profile from a 4 : 1 SR-28 form during a column leaching test using a 7 ml min-1 flow rate.

The greatest peak and stabilized fluxes occurred in SR-HP forms 19-28 (Table 3-3)

(Figures 3-5a-l). Holes were drilled into these forms to increase the surface area, leading

-1 to greater peak H2O2 fluxes on average (36.6 mg min ), compared to forms without holes

(7.1 mg min-1), or 518% greater. Drilling lateral and longitudinal holes into these forms increased surface area by 39.9% on average (s= 0.06, n=10). The final 6 SR-HP forms were manufactured with the same, higher, oxidant : resin ratio (4:1), with similar hole patterns to increase surface area, leading to higher peak and stabilized fluxes compared to

SR-HP 19-22, made with 3.5:1 and 3.75:1 mixing ratios and only longitudinal hole 60 patterns (42.2 mg min-1 and 1.1 mg min-1 compared to 36.6 and 0.8 mg min-1) (Figures 3-

5g-l).

Concentration measurements were recorded at the same times for SR-HP 23-28 over a two-week span. Figure 3-6 displays the average flux during this time with one standard deviation error represented for each sampling point on a semi-log (Fig. 3-6a) and normal scale (Fig. 3-6b). The greatest amount of variance occurs during the peak release, ranging from 29.9 – 52.1 mg min-1. Hydrogen peroxide fluxes from the 6 SR-HP column tests are included in Appendix A.

40 ) 1 -

(mg (mg min

20 O 2 H

0 0 100 200 300 400 Time (hr)

Figure 3-6a. Average release profile of SR-HP 23-28 with one standard deviation error bars.

61 60

40 ) 1 -

(mg min

20 O 2 H

0 0.1 1 10 100 1000 Time (hr)

Figure 3-6b. Average release profile of SR-HP 23-28 with one standard deviation error bars.

Stabilized release rates (after 5 hours) were modeled with first order decay equations to compare the consistency release profiles among SR-HP 23-28. Two-sided t-

Tests performed between constants obtained from the modeled equations and standard error suggests variability among these forms. P-values range between <0.001 and 0.51, both rejecting and failing to reject a null hypothesis of differing stabilized release rates among SR-HP 23-28 (Table 3-4). Slow-release form 26 showed the greatest variability from the other forms, producing statistically significant p-values in each case. A 0.007 p- value from T-tests performed on the average stabilized flux between SR-HP 23-28 and 62 SR-HP 19-22 (made with oxidant : resin < 4) implies different release profiles are obtained with forms made from different mixing ratios and surface areas.

Table 3-4. P-values from two-sided T-tests performed between first order decay constants between SR-HP 23-28 SR#HP& 24& 25& 26& 27& 28& 23& 0.057% 0.002% 0.000% 0.003% 0.359% 24& % 0.154% 0.000% 0.305% 0.491% 25& % % 0.000% 0.515% 0.073% 26& % % % 0.000% 0.000% 27& % % % % 0.131%

3.2.3 Oxidant Release Efficiency / Recovery Rate

Increasing the oxidant : resin mixing ratio tends to increase the peak release rate; however, column leaching test results indicate ratios greater than 4 : 1 generally yield decreasing peak release rates (Figure 3-7). A ratio of 4 : 1 was found to be ideal for both peak and stabilized release rates. These forms had the structural integrity to release H2O2 for an extended amount of time (> 2 weeks). Drilling holes into the forms, exposing more area for diffusion to occur had the greatest impact on peak and stabilized release efficiency (Figure 3-8). Peak release rate plotted against surface area generates a Pearson correlation coefficient of R=0.91. 63 60

) 1 - ]min 2 O

2 40

20 Peak Release (mg[H Release Peak 0 0 2 4 6 8 Oxidant : Resin Mixing Ratio

Figure 3-7. Peak H2O2 flux versus oxidant : resin mixing ratio for all SR-HP forms

60 ) 1 - ]min 2 O 2 40

Peak Release (mg[H Peak 0 0 50 100 150 200 250 Surface Area (cm2)

2 Figure 3-8. Peak H2O2 flux versus surface area (cm ) for all SR-HP forms

64 3.3 Proof-of-Concept Iron Removal Tests

Experiments were preformed in the laboratory to simulate Fe2+ removal with SR-

HP forms using Fenton’s reagent. Nineteen tests were conducted using different

+2 concentrations of both Fe and H2O2 to measure the removal efficiency for different

+2 [Fe ]/[H2O2] ratios. Water collected from nearby AMD sites were used in these experiments. Concentrated Fe2+ solutions were added to these waters to obtain a range of

Fe2+ to test for removal. In keeping with the results of Mahiroglu et al., (2009), initial

+2 [Fe ]/[H2O2] ratios approached 1.8 for optimal removal efficiency. Final ratios of

+2 [Fe ]/[H2O2] ranged from 0.5 to 10.5 for all 19 tests. Appendix B contains the weight of

+2 2+ sodium percarbonate (kg) in one day necessary for a [Fe ]/[H2O2] 2:1 for different Fe concentrations (mg L-1) and discharges (L s-1).

2+ The removal of Fe by H2O2 was rapid in all proof-of-concept tests (Figure 9a-c), oxidizing a majority of the Fe+2within 1 minute of mixing. Experiments with

+2 2+ [Fe ]/[H2O2] < 2 were able to oxidize over 90% of Fe after 10 minutes (Figure 3-9c).

+2 2+ Solutions with greater [Fe ]/[H2O2] ratios (2-10) were less efficient at removing Fe , but still able to oxidize 60-80% of the available Fe+2. Conductivity and pH were monitored before the solutions were mixed and 10 minutes after (Table 3-5).

Conductivity and pH ranged between 1947-2285 and 2.03-2.18 respectfully before the two solutions were mixed. After treatment conductivity and pH values ranged between

215-941 and 2.35-3.22 respectfully. Conductivity was reduced by 69% and pH levels rose 0.61 on average. Results from the proof-of-concept tests helped determine the amount of SR-HP forms to be installed at the field study site. 65 Table 3-5. pH and Conductivity before and 10 min after proof-of-concept iron removal tests for different [Fe+2]/[H2O2] ratios Conductivity* Conductivity* [Fe+2]/[H2O2]* pH*0*min* pH*10*min* (μS/cm)* (μS/cm)* *0*min* *10*min* 0.50* 2.12* 3.22* 1787* 465* 0.91* 2.18* 2.71* 1450* 941* 0.97* 2.12* 2.6* 1787* 595* 1.50* 2.02* 2.61* 1825* 592* 1.74* 2.01* 2.67* 1970* 464* 1.80* 2.02* 2.6* 1825* 477* 1.80* 1.95* 2.83* 2284* 338* 1.82* 2.08* 2.55* 1862* 758* 1.85* 2.01* 2.57* 1970* 487* 2.04* 2.02* 2.65* 1927* 588* 2.10* 1.95* 3.08* 2285* 215* 3.02* 2.02* 2.62* 1922* 611* 4.23* 2.02* 2.6* 1922* 621* 4.87* 2.02* 2.45* 2197* 580* 5.90* 2.02* 2.45* 2197* 638* 7.25* 2.02* 2.35* 1944* 752* 10.50* 2.00* 2.39* 1944* 718* Ave$ 2.03$ 2.64$ 1947$ 579$ Min$ 1.95$ 2.35$ 1450$ 215$ Max$ 2.18$ 3.22$ 2285$ 941$ SD$ 0.06$ 0.22$ 208$ 169$ 66 100%

75% Removed

2+ 50%

25% Percentage of Fe Percentage 0% 0 3 6 9 12 2+ [Fe ]/[H2O2]

+2 Figure 3-9a. Ferrous iron removal efficiencies after 1 minute at different [Fe ]/[H2O2] ratios

100%

75% Removed

2+ 50%

25%

Percentage of Fe Percentage 0% 0 3 6 9 12 [Fe 2+]/[H O ] 2 2 +2 Figure 3-9b. Ferrous iron removal efficiencies after 5 minutes at different [Fe ]/[H2O2] ratios

67 100%

75%

50%

25%

0% 0 3 6 9 12 II [Fe ]/[H2O2]

+2 Figure 3-9c. Ferrous iron removal efficiencies after 10 minutes at different [Fe ]/[H2O2] ratios

3.4 Small-Scale Field Application Test

Fourteen SR-HP forms were placed in a netted sack and anchored upstream of

AMD seep BH00690 (1 in Figure 3-10)(Table3-5). If these SR-HP forms obtained a peak flux similar to the average peak fluxes obtained from SR-HP 23-28 (~42.2 mg min-1),

-1 +2 st with a flow of 0.7 Ls , this would create a [Fe ]/[H2O2] of ~3 within the 1 hour of treatment. Based on results from laboratory column leaching and proof-of-concept

2+ removal tests, this provides enough H2O2 to oxidize 70-80% of Fe available from

BH00690 at the peak release (14*42.2= 592 mg min-1), 1 hour after treatment. After 1 day, H2O2 fluxes from SR-HP forms decreased exponentially, failing to provide ample amounts of oxidants (Figure 3-21). The six established sampling locations were 68 2+ monitored over 2 weeks for H2O2 and Fe concentrations, field water chemistry, and discharge (Figures 3-10 - 3-21 and Table 3-7).

Each time the field site was visited one SR-HP form was removed, dried, and weighed to determine how much Na2CO3·1.5H2O2 had dissolved. These dried SR-HP forms were placed into glass columns and performed leaching experiments to determine how much, if any, H2O2 remained present in the SR-HP forms (Table 3-6). In all cases,

-1 negligible H2O2 concentrations (<1 mg L ) were obtained upon column leaching tests after five hours. Mass differences calculated in Table 3-6 suggest H2O2 remained in SR-

HP 5, 6, 8, and 13; however, these forms did not generate sufficient diffusion pathways and the outer coating of resin prohibited further dissolution (Figure 3-11).

Table 3-5: Properties of SR-HP forms used in field removal demonstration SR Mass Diameter Height Volume Density Mass Mass 3 -3 Form Ratio (g) (cm) (cm) (cm ) (gcm ) Oxidant (g) H2O2 (g) 1 4:1 113.3 2.8 12.8 78.8 1.43 90.64 29.5 2 4:1 91.1 2.8 11.2 69.0 1.32 72.8 23.7 3 4:1 102.9 2.8 12.7 78.2 1.32 82.3 26.8 4 4:1 123.6 2.8 13.7 84.4 1.47 98.9 32.1 5 4:1 109.1 2.8 12.7 78.2 1.40 87.3 28.4 6 4:1 107.2 2.8 13.0 80.0 1.34 85.8 27.9 7 4:1 86.8 2.8 10.1 62.2 1.40 69.4 22.6 8 4:1 100.8 2.8 11.8 72.7 1.39 80.6 26.2 9 4:1 107.2 2.8 12.4 76.4 1.40 85.8 27.9 10 4:1 119.5 2.8 13.6 83.7 1.43 95.6 31.1 11 4:1 103.1 2.8 11.3 69.6 1.48 82.5 26.8 12 4:1 71.5 2.8 8.8 54.2 1.32 57.2 18.6 13 4:1 84.4 2.8 9.8 60.3 1.40 67.5 21.9 14 4:1 82.7 2.8 9.6 59.1 1.40 66.2 21.6

69 Table 3-6: Percentage of dissolved oxidant from SR-HP forms removed from field demonstration SR Form Mass (g) End Mass (g) Time Removed (hr) Recovery Rate (%) 5 109.1 89.44 52 22.5 6 107.2 76.18 99 37.8 8 100.8 58.6 24 52.3 13 84.4 32.94 146 76.2 14 82.7 11.15 190 100

Table 3-7: Discharge measured during the small-scale field demonstration at BH00690 and upstream and downstream measurements Upstream BH00690 Downstream Time(hr) (L s-1) (L s-1) (L s-1) 0.00 4.26 1.07 5.33 0.08 4.26 1.07 5.33 0.25 4.26 1.07 5.33 0.50 4.26 1.07 5.33 0.75 4.26 1.07 5.33 1.00 4.26 1.07 5.33 1.50 4.26 1.07 5.33 2.00 4.26 1.07 5.33 4.00 4.26 1.07 5.33 4.50 4.26 1.07 5.33 5.00 4.26 1.07 5.33 23.00 3.69 1.15 4.84 23.50 3.69 1.15 4.84 24.00 3.69 1.15 4.84 52.00 6.84 1.18 4.40 99.00 3.21 0.96 6.84 146.00 3.21 0.92 5.79 190.00 3.21 0.98 7.39 Min 3.21 0.92 4.40 Max 6.84 1.18 7.39 Average 4.13 1.07 5.42 SD 0.79 0.06 0.69

Figure 3-10. Sketch map of BH006900 in relation to Brush Fork and sampling sites.

Figure 3-11. Dissolution of SR-HP forms varies depending on the heterogeneous nature of mixing. SR-HP forms with little resin on exterior exhibit much quicker dissolution.

Manufacturing SR-HP forms with similar dissolution tendencies and consistent release rates continues to be a challenge. All 14 SR-HP forms were produced with the same 4:1 oxidant : resin; however, after 3 days these forms were not experiencing similar diffusion.

Ferrous iron concentrations decreased after installation of SR-HP forms (Figures

3-12 and 3-19 - 20). The greatest removal (81%) was obtained after one hour, from 18.7 to 3.5 mg L-1 Fe2+; however, background sampling of the 5 sites along Brush Fork downstream of BH006900 suggest much of the iron oxidizes immediately after exiting the pipe, cascading down 1.5 meters before entering Brush Fork. In addition, after 24

2+ hours of treatment, H2O2 concentrations fell below the detection limit, yet Fe concentrations measured were 70% less than the background levels, suggesting 72 significant natural variability in Fe2+ concentrations along the sampled section of Brush

Fork.

100%

80%

60% removed 2+ 2+ 40%

20% Percent Fe Percent 0% 0 50 100 150 200 Time (hr)

Figure 3-12a . Percent of Fe2+ removed during the field demonstration (normal time scale).

100%

80%

60% removed 2+ 2+ 40%

20% Percent Fe Percent 0% 0.01 0.1 1 10 100 1000 Time (hr)

Figure 3-12b. Percent of Fe2+ removed during the field demonstration (log time scale). 73 pH levels rose three points 5 hours after the addition of SR-HP forms (Figure 3-15) with the greatest difference occurring at sampling location 2, the closest measuring point to BH00690 (Figure 3-16). By the end of the field demonstration pH values had returned to background levels. Conductivity changes measured at the six sampling locations were also greatest at location 2, decreasing by 54% within 45 minutes of treatment from SR-

HP forms (Figures 3-17 and 3-18).

Hydrogen peroxide concentrations measured immediately upstream of the confluence with BH00690 followed similar release patterns to laboratory column leaching experiments (Figure 3-21). The measured peak flux (706 mg min-1) exceeded the estimated peak flux (592 mg min-1) by 19%, perhaps due to the stronger flow the SR-

HP forms were subjected to in the field (4.26 L s-1) compared to laboratory column leaching experiments (7-8 mL min-1). Six of the SR-HP forms used in the field demonstration did not remain structurally intact throughout the test and were found in pieces after the first day. The stronger flow likely dissolved and broke apart these forms more easily compared to the weaker flows used in the laboratory column leaching experiments, suggesting a lower oxidant : resin mixing ratio or stronger binder may be desired for longer field application.

Figure 3-13. Site BH00690 looking upstream, location of SR-HP treatment forms.

Figure 3-14. Iron precipitating after one day of treatment at sample site #3 (24 meters downstream of mixing point. 75 9 8 7 6

4 pH 3 2 2 3 4 1 5 6 1 0 0 100 200 300 400 Time (hr) Figure 3-15a. pH versus time measured at six different sampling locations (normal time scale).

9 8 7 6

4 pH 3 2 2 3 4 5 6 1 1 0 0.01 0.1 1 10 100 1000 Time (hr) Figure 3-15b. pH versus time measured at six different sampling locations (log time scale).

pH 4

0 0.25 1 4 2

24 99 190 336 1

0 -20 -10 0 10 20 30 40 50 Distance from BH00690 (m)

Figure 3-16. pH versus distance from BH006900 measured at different times.

77 700

600

500

400 (μS/cm)

300

200

2 3 4 Conductivity 100 5 6 1 0 0 100 200 300 400 Time (hr) Figure 3-17a. Conductivity versus time measured at the six different sampling locations.

700 600

500 400 (μS/cm)

300

200 2 3 4 100 Conductivity Conductivity 5 6 1 0 0.01 0.1 1 10 100 1000 Time (hr)

Figure 3-17b. Conductivity versus time measured at the six different sampling locations.

700

600

500

400

300 Conductivity (μS/cm) Conductivity 200 0 0.25 1 4 100 24 99 190 336 0 -20 -10 0 10 20 30 40 50 Distance from BH00690 (m)

Figure 3-18. Conductivity versus distance from BH006900 measured at different times.

79 25 2 3 4 5 6 20

15 (mg/L)

10 2+ Fe

0 0 100 200 300 400 Time (hr)

Figure 3-19a . Fe2+ versus time measured at five sampling locations downstream of BH006900.

25 2 3 4 5 6 20

15 (mg/L)

10 2+ Fe

0 0.01 0.1 1 10 100 1000

Time (hr) Figure 3-19b. Fe2+ versus time measured at five sampling locations downstream of BH006900. 80

0 25 0.25

1 20 4

24 15 99 (mg/L)

190 2+

10 Fe 336

0 0 10 20 30 40 50 Distance from BH00690 (m)

Figure 3-20. Fe2+ versus distance from BH006900 measured at different times.

81 3

2.5

2 - mgL 1.5 2 O 2

1 H

0.5

0 0.01 0.1 1 10 100 1000 Time (hr)

Figure 3-21. H2O2 versus time measured immediately upstream of BH006900.

10000

1000 ) 1 -

100

10 (mgmin Flux Fe(II) H2O2 1 0 50 100 150 200 Time (hr) 2+ Figure 3-22a. Fe measured at sampling location 2 and H2O2 fluxes versus time. 82 10000

1000

) 1 -

100 Flux (mgmin Flux Fe(II) 10

1 0.01 0.1 1 10 100 1000 Time (hr)

2+ Figure 3-22b. Fe measured at sampling location 2 and H2O2 fluxes versus time.

] 2 O

30 2 ]/[H

20 2+ [Fe

0 0.01 0.1 1 10 100 Time (hr) 2+ Figure 3-23. [Fe ]/[H2O2] ratio versus time.

83 2+ Optimal [Fe ]/[H2O2] ratios were measured within the first hour of the field demonstration (Figure 3-23). Iron concentrations measured at sampling site 2 were used in this calculation, the location closest to BH00690. During the first two hours of the

2+ field demonstration [Fe ]/[H2O2] ratios ranged between 0.73 – 4.60. While 70 to 100 % of Fe2+ was expected to be removed with Fenton’s Reagent based on results in Figure 3-

9a-c, 56 – 81 % of Fe2+ was removed from the field demonstration (Figures 3-12 and 3-

2+ 23). After 4 hours the [Fe ]/[H2O2] ratio increased greatly, suggesting oxidation from

H2O2 was not as effective. These results suggest feasibility of using SR-HP to treat oxidizable metals in AMD water. Further development of SR-HP forms with higher release rates, longer durations, stronger binder, and improved reproducibility is warranted.

84 CHAPTER 4: CONCLUSION

This study focused on developing SR forms that release both H2O2 and carbonates at predetermined rates over an extended amount of time to oxidize Fe2+ and increase alkalinity. A combination of laboratory column leaching tests, proof-of-concept Fe2+ removal tests, and chemical sampling at BH00690 helped determine the correct number of SR forms for optimal treatment. Column leaching tests confirmed an ideal oxidant : resin mixing ratio of 3.8:1 – 4.2:1. Slow-release forms made with this mixing ratio generated the highest average peak release rates (22.5 mg min-1) followed by higher sustained release rates (5.2 x 10-1 mg min-1), compared to forms made outside of this oxidant : resin mixing ratio (8.8 mg min-1 and 2.7 x 10-1 mg min-1 respectfully.

Additionally these forms produced higher recovery rates (52.4%) compared to other forms (28.7%) and produced sustained release rates for 14 days compared to 7 days.

Increasing the surface area of SR forms had a great impact on the peak and sustained release rates. Latitudinal and longitudinal holes drilled into the final nine forms

(SR 23-28) increased the surface area of these forms by 39.9% on average. The peak release rates from these forms were on average 518% greater than SR forms without holes (36.6 mg min-1 compared to 7.1 mg min-1). Statistical analyses were preformed on release rates of SR 23-28 to determine consistency (Figure 3-30 and Table 3-4). The greatest variability occurs during the peak release (hours 0-8) followed by more consistent, slower release rates.

Laboratory proof-of-concept Fe2+ removal tests helped determine the most

+2 efficient [Fe ]/[H2O2] ratios for treatment. Results from these tests indicated 85 +2 2+ +2 [Fe ]/[H2O2] ratios <2 were able to oxidize 90% of Fe , while [Fe ]/[H2O2] ratios were able to remove only 60-80% of Fe2+ over the same time.

Baseline chemical sampling results at BH00690 and the five downstream field- sampling sites indicate that much of the Fe2+ is removed naturally from as the water drains from BH00690, falling approximately four feet to the stream level. Ferrous iron concentrations measured at BH00690 (118 mg L-1) were almost ten times greater than

Fe2+ concentrations measured in Brush Fork prior to treatment (19 mg L-1). Ferrous iron concentrations reductions were greatest after one hour of treatment (81% reduction).

After the first day of treatment, H2O2 concentrations were measured below the detection limit; however, Fe2+ were measured at 73% less of the values prior to treatment, suggesting natural variability in Fe2+ along this section of Brush Fork in addition to dilution from the clean, upstream source. Additional sampling along this section of stream is recommended for future analysis of Fe2+ variability for optimal treatment.

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90 APPENDIX A: H2O2 FLUXES FROM 6 SR-HP COLUMN TESTS OVER TWO

WEEKS

Time (hr) SR-HP Form H O Flux (mg min-1) 2 2 Standard Average 0 23 24 25 26 27 28 Deviation 0.25 7.46 13.92 10.99 3.27 14.75 14.58 4.64 10.83 0.5 28.84 29.32 29.16 29.62 29.32 29.33 0.26 29.27 0.75 36.40 47.72 49.78 52.13 29.93 37.55 8.86 42.25 1.25 25.39 41.39 32.59 39.01 18.45 24.45 8.98 30.21 1.75 16.54 30.10 24.56 28.22 7.98 21.51 8.20 21.48 3 7.74 12.22 13.24 18.87 3.95 10.57 5.07 11.10 4.5 3.92 6.59 5.65 9.57 2.36 5.34 2.46 5.57 8 2.61 3.94 3.09 5.05 1.94 3.26 1.08 3.31 11 2.65 2.33 2.07 5.54 1.99 2.24 1.36 2.80 24 2.05 1.76 1.59 4.29 1.52 1.69 1.06 2.15 30 1.93 1.28 1.32 4.23 1.40 1.23 1.17 1.90 48 1.89 1.07 1.02 2.11 1.14 0.59 0.58 1.30 54 1.75 1.20 0.88 1.85 1.28 0.80 0.44 1.29 75 1.46 1.05 1.21 0.54 1.13 0.80 0.32 1.03 120 0.89 0.84 0.73 0.43 0.75 0.38 0.21 0.67 144 0.31 0.22 0.45 0.50 0.27 0.23 0.12 0.33 169 0.19 0.27 0.48 0.41 0.20 0.29 0.12 0.31 192 0.26 0.21 0.37 0.27 0.24 0.20 0.06 0.26 337 0.10 0.13 0.26 0.17 0.18 0.15 0.05 0.16

91 APPENDIX B: WEIGHT OF SODIUM PERCARBONATE (KG/DAY) NECESSARY

+2 2+ -1 FOR A [FE ]/[H2O2] 2:1 FOR DIFFERENT FE CONCENTRATIONS (MG L ) AND

DISCHARGES (L S-1)

Fe2+ Discharge L s-1 (mg L-1) 0.25 0.50 0.75 1.00 1.25 1.50 1.75 2.00 5 101 202 303 405 506 607 708 809 10 202 405 607 809 1012 1214 1416 1618 15 303 607 910 1214 1517 1821 2124 2428 20 405 809 1214 1618 2023 2428 2832 3237 25 506 1012 1517 2023 2529 3035 3540 4046 30 607 1214 1821 2428 3035 3641 4248 4855 35 708 1416 2124 2832 3540 4248 4956 5664 40 809 1618 2428 3237 4046 4855 5664 6474 45 910 1821 2731 3641 4552 5462 6372 7283 50 1012 2023 3035 4046 5058 6069 7081 8092 55 1113 2225 3338 4451 5563 6676 7789 8901 60 1214 2428 3641 4855 6069 7283 8497 9710 65 1315 2630 3945 5260 6575 7890 9205 10520 70 1416 2832 4248 5664 7081 8497 9913 11329 75 1517 3035 4552 6069 7586 9104 10621 12138 80 1618 3237 4855 6474 8092 9710 11329 12947 85 1720 3439 5159 6878 8598 10317 12037 13756 90 1821 3641 5462 7283 9104 10924 12745 14566 95 1922 3844 5766 7687 9609 11531 13453 15375 100 2023 4046 6069 8092 10115 12138 14161 16184

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