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Arsenic removal by iron oxides Aredes, Sonia 2005

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ARSENIC REMOVAL BY IRON OXIDES BY  SONIA AREDES  B.A.Sc. Chemical Engineering. Universidad Nacional de Salta. Argentina. 1981  A THESIS SUBMITTED IN PARTIAL F U L F I L L M E N T OF REQUIREMENTS FOR THE DEGREE OF M A S T E R OF APPLIED SCIENCE In  THE F A C U L T Y OF G R A D U A T E STUDENTS (Mining Engineering)  THE UNIVERSITY OF BRITISH C O L U M B I A December, 2005  ©Sonia Aredes, 2005  ABSTRACT  Arsenic has long been recognized as a toxin and carcinogen. Arsenic contaminated drinking water probably poses the greatest threat to human health. A successful treatment for removing arsenic from drinking water requires an understanding of arsenic chemistry and the physical-chemical processes that occur during each treatment step. Iron oxide (Fe-Ox) minerals showed good efficiency for arsenic removal (Simeonova 2000, Matis et.al., 1999). Additionally, naturally occurring iron oxides are more attractive for arsenic removal from contaminated water than the synthetic oxides because they are more cost effective. However, few studies have been carried out on the feasibility of their use as adsorbents for arsenic removal. Hematite, magnetite, goethite and laterite have been studied in their role as arsenic adsorbents. Results showed that all of them are suitable as arsenic adsorbents. Electroacoustic Tests (ET) tests showed that arsenic adsorption occurs over the whole pH range considered (4-11) and also that the Fe-Ox have IEP at pH between 6.5 and 8.5. Their surface charge is negative at pH<pzc and positive at pH >pzc. The presence of inner sphere complex, which implies stability of the arsenic adsorbed onto Fe-Ox because of covalent bonding, was shown by ET and leaching tests. Leaching tests by MgCL; were performed to study the stability of the adsorption products and results expressed on a weight percentage basis showed that hematite had 60.2%, magnetite 75.4%, goethite 78.0% and laterite 86.2% of arsenic strongly fixed. While these results expressed on a 2  *  2  surface area basis showed that hematite had 0.16mg/m , magnetite O.llmg/m , goethite 0.065 mg/m and laterite 0.011 mg/m of arsenic strongly fixed.  ii  In addition, this study presents a simple method developed to remove arsenic from water using natural iron oxides (Fe-Ox) minerals. The method involves mixing natural iron mineral bearing soils (lateritic soils) with arsenic contaminated water for ten minutes and then filtering (coffee filter). The aadsorption capacity of laterite was estimated at 0.1 lmg/m . After addition, arsenic levels in the treated water were below drinking water standards. The treatment method is inexpensive and simple, making it suitable for house hold use.  iii  TABLE OF CONTENTS  ABSTRACT  ii  TABLE OF CONTENTS  iv  LIST OF TABLES  viii  LIST OF FIGURES  x  ACKNOWLEDGEMENTS  xii  1. INTRODUCTION  1  1.1 STATEMENT OF THE PROBLEM  1  1.2 THESIS OBJECTIVE AND SIGNIFICANCE  4  1.2.1 Objective  4  1.2.2 Significance and Contribution of the Work  4  2. LITERATURE REVIEW  6  2.1 ARSENIC IN THE ENVIRONMENT  6  2.1.1 Arsenic in Drinking Water  6 National and International Standards for Drinking Water  6 World Distribution of Groundwater Arsenic Problems  6 Arsenic in B. C.  7 Health Effects Associated with Arsenic in Drinking Water 2.1.2 Sources of Arsenic in the Environment  8 11 Natural Sources of Arsenic  11 Anthropogenic Sources of Arsenic  13  iv  2.1.3 Environmental Transfer of Arsenic  14 Arsenic Transport between Media  14 Arsenic in Soils  16 Arsenic in Water  18 Arsenic in Sediments  22 The Role of Bacteria in Arsenic Mobilization  26  2.1.4 Arsenic Removal Summary of Existing Technologies  27 Precipitation - Coprecipitation  28 Membrane Filtration  28 Adsorption Treatment  29 Ion Exchange  29  2.2 IRON OXIDES  29  2.2.1 Natural Iron Oxide Occurrence and Description  29  2.2 A A Introduction  29 Iron Oxides in Rocks and Ores Iron Oxides in Soils  31 35  2.2.2 Iron - Arsenic Compounds Interaction in the Environment  36  2.2.3 The Solubility of Fe-Oxides  37  2.2.4 The Adsorption Process and Iron Oxides  39 Surface Chemistry  39  2.2.4. 2 Arsenic (V) Adsorption onto Iron Oxides 2.2.5 Conclusions  44 48  v  3. EXPERIMENTAL PROCEDURES  49  3.1 MATERIALS  49  3.1.1 Solids  49  3.1.2 Solutions  49  3.2 PROCEDURES  50  3.2.1 Preliminary Tests  50  3.2.2 Scoping Tests  50  3.2.3 Electroacoustic Tests  52  3.2.4 Adsorption & Leaching Tests  56  3.2.5 Adsorption Isotherms  60  3.2.6 Evaluation of a Simple Water Treatment Process  62  4. RESULTS AND DISCUSSION  63  4.1 SCOPING TESTS  63  4.2 ELECTROACOUSTIC TESTS  64  4.2.1 Magnetite  64  4.2.2 Hematite  65  4.2.3 Goethite  69  4.2.4 Laterite  70  4.3 ADSORPTION & LEACHING TESTS  72  4.3.1 Adsorption Test  72  4.3.2 Leaching Test  72  4.4 ADSORPTION ISOTHERMS  75  4.5 EVALUATION OF A SIMPLE WATER TREATMENT PROCESS  78  vi  4.6 DISCUSSION OF RESULTS 4.6.1 Electro acoustic Measurements 4.6.2 Adsorption Tests 4.6.3 Leaching Test 4.6.4 Adsorption Isotherms 4.6.5 Evaluation of Simple Water Treatment Process 4.6.6 Comparison between Fe-Ox  5. CONCLUSIONS & RECOMMENDATIONS 5.1 CONCLUSIONS 5.2 RECOMMENDATIONS REFERENCES APPENDIX  LIST OF TABLES  Table 2.1 Arsenic Regulatory Limits  6  Table 2.2 Arsenic Concentrations in Environmental media  11  Table 2.3 Most Common Minerals that Contain Arsenic  12  Table 2.4 Summary of Current and Past Uses of Arsenic  13  Table 2.5 Most Common Iron Oxides  30  Table 2.6 Most Common Iron Oxides-Hydroxides and Hydroxides  30  Table 2.7 Dominant Occurrences of the Fe-Ox in Geological Formations  31  Table 2.8 Summary of the Occurrence of Iron Oxides in Various Soils  35  Table 3.1 Solid Particle Size & Surface Area Analysis  50  Table 3.2 Mineralogical Composition of Samples  51  Table 3.3 Test Conditions for Hematite Titrated at Constant pH with As (V) Solution  55  Table 3.4 Test Conditions for Titration of Fe-Ox Slurries in the Whole pH Range 55 Table 3.5 Tests Conditions for Titration of Fe-Ox Slurries Loaded with As (V)  55  Table 3.6 Sequential Extraction Conditions  57  Table 3.7 Adsorption Tests at Constant pH = 5.2  60  Table 3.8 Adsorption Tests at Constant pH = 9.2  61  Table 3.9 Test Conditions for the Evaluation of a Simple Water Treatment Process  62  Vlll  Table 4.1 Scoping Tests. As Adsorbed from Different Fe-Ox Slurries at pH 5.0  63  Table 4.2 Adsorption Test Results  73  Table 4.3 Leaching Test Results  75  Table 4.4 Leaching Test Results in a Surface Area Basis  75  Table 4.5 Adsorption Isotherm at pH 5.2  76  Table 4.6 Adsorption Isotherm at pH 9.2  77  Table 4.7 Evaluation of a Simple Water Treatment Process  79  ix  LIST OF FIGURES Figure 2.1 Distributions of Documented World Problems with As in Groundwater 7 Figure 2.2 Sources of Human Exposure to Arsenic and Various Modes of Arsenic Toxicity  10  Figure 2.3 As Cycle. Environmental Transfer of Arsenic  15  Figure 2.4 Chemical Forms of Arsenic and their Transformation in Soils  17  Figure 2.5 Eh- pH Diagram for the System As-O-H-S  20  Figure 2.6 Eh - pH Diagram for the Systme As - S - Fe - H2O  37  Figure 2.7 Solubilities of Goethite, Ferrihydrite and "Soil-Fe" as a Function of pH 38 Figure 2.8 Surface Complex Formation of an Ion (e.g. Cation) on the Hydrous Oxide Surface  42  Figure 3.1 Sequential Extraction Procedure  59  Figure 4.1 ESA vs. pH of Magnetite Slurries  65  Figure 4.2 ESA vs. pH of Hematite Slurries  66  Figure 4.3 ESA vs. pH of Hematite Slurries, Back and Forward Titration  67  Figure 4.4 Titration of Hematite Slurry with Arsenic (V) Solution  68  Figure 4.5 ESA vs. pH of Goethite Slurries  69  Figure 4.6 ESA vs. pH of Laterite#l Slurries  70  Figure 4.7 ESA vs. pH of Laterite#2 Slurries  71  Figure 4.8 Adsorption Isotherm at pH 5.2  76  Figure 4.9 Adsorption Isotherm at pH 9.2 for Laterite #2  77  Figure 4.10 Langmuir Isotherm Model at pH 9.2  78  Figure 4.11 As Fixed (%) vs. Fe-Ox Surface Area  85  Figure 4.12 As Adsorbed & ESA vs. As Initial  87  xi  ACKNOWLEDGMENTS This thesis is a product of collaboration and support of many people, to all of them thanks. My deepest gratitude to Dr Bern Klein, who with constructive criticism, leads me throughout all the stages of my thesis. He was always there when I needed his advice, and watched my work objectively to advise me on the best way to follow. I consider that his support was fundamental for the completion of my thesis. I want to also thank Dr Marek Pawlik, who leads me throughout the knowledge of the surface chemistry; I enjoyed and learnt from his lectures, he had always a good disposition to listen to me. Many thanks to Dr Marcello Veiga, for encouraging me to pursue this degree in the UBC Mining department, and for providing valuable thoughts and observations. I want to thank in a special way for the support received from Sally Finora, Leslie Lerox, Pius Lo, Maria Lui, and Keith Jellis, they always assisted me with their best predisposition, facilitating my tasks. Finally, to my wonderful friends in the Mining department, thanks to their help and support.  xii  1. INTRODUCTION  1.1 STATEMENT OF THE PROBLEM Arsenic is a common element found in the atmosphere, soils, rocks, natural waters and organisms. It is mobilized in the environment due to a combination of natural processes such as weathering reactions, biological activity, volcanic emission, and through a range of anthropogenic activities. The majority of environmental arsenic problems are thus a consequence of mobilization under natural conditions, but anthropogenic activities have had an important impact through activities such as mining, combustion of fossil fuels, the use of arsenic-based pesticides, herbicides and crop desiccants. Arsenic has long been recognized as a toxin and carcinogen. Chronic exposure to arsenic can occur through various sources, both natural (e.g. weathering reactions, volcanic emissions) and man made (e.g. mining related activities). Of the various sources of arsenic in the environment, exposure to arsenic through contaminated drinking water probably posses the greatest threat to human health. Drinking water comes from a variety of sources depending on local availability: surface water (rivers, lakes, reservoirs and ponds), groundwater (aquifers) and rain water. In most of the cases, high concentrations of arsenic are mainly found in ground water. Wellknown high-As groundwater areas have been found in Argentina, Chile, Mexico, China, Hungary, West Bengal (India), Bangladesh and Vietnam. Contamination of groundwater by arsenic in Bangladesh is the largest poisoning of a population in history, with millions of people exposed (Smith et al, 2000).  1  Epidemiological studies implicate arsenic as a cause of diseases such as hyperkeratosis, peripheral vascular disease ("black foot disease"), skin, lung and bladder cancers, diabetes, anemia, and disorders of the immune nervous and reproductive systems. The USEPA states "Arsenic is the only carcinogen where exposure through drinking water has been clearly demonstrated to cause human cancer" (Murphy and Guo, 2003). Most of the countries have set arsenic limits for drinking water at 50 ppb. However, following the accumulation of evidence for the chronic toxicological effects of As in drinking water, regulatory limits tend to be reduced from this old standard. •  The Canadian Federal-Provincial-Territorial Committee on Drinking Water set the standard for arsenic in drinking water at 25 ppb as a Maximum Acceptable Concentration (MAC) in April 2004  •  On January 22, 2001 the USEPA adopted a new standard for arsenic in drinking water at 10 ppb, replacing the old standard of 50 ppb.  Most of arsenic removal technologies are referred to as "conventional water treatment". Frequently, they are either inadequate or too expensive to be suitable for the small communities in developing countries. In addition, technology selection for the arsenic treatment of drinking water has to take into account the complex chemistry of arsenic that may affect water treatment. Most available technologies are inappropriate for rural communities in developing countries because of the following: •  They are mainly designed for centralized systems (these rural communities usually do not have centralized water treatment facilities)  2  •  They require a given system size, and require water of certain quality (i.e. pH range)  •  Require skilled operators  •  Generate waste that needs to be treated further, and  •  They are excessively expensive for small villages of developing countries.  Therefore, a simple and affordable treatment method is needed which will not require water pre-treatment (i.e. pH adjustment) and will satisfy water needs for drinking and food preparation (only 2 - 3 liters per capita per day). Arsenic is commonly found in water as inorganic species. As (V) or As (III) occur as arsenate  (H2ASO4"  and  HA.SO4 ") 2  and arsenite (H3ASO3) ions which are the dominant  species in water. Some conditions that affect arsenic speciation in water include: pH, redox potential, the presence of other ions, and microbial activity. The unstable nature of arsenic species makes it difficult to treat, resulting in treated wastes whose toxicity and mobility can change with environmental conditions. Therefore, the successful treatment and long term disposal of arsenic wastes requires an understanding of arsenic chemistry and the physical-chemical processes that occur during each treatment step. The use of naturally occurring iron oxides for arsenic removal from water is promising. They are abundant in nature and thus cheap, though few studies have been carried out on the feasibility of their use as adsorbents for arsenic removal. This study shows the suitability of using different natural iron oxides for arsenic removal. Arsenic-iron oxide interactions were studied through electroacoustic (surface charge), adsorption and leaching tests. Also, the suitability of using a lateritic soil for arsenic removal from water is demonstrated. Thus, arsenic removal in rural areas can be carried  3  out at the household level. The system proposed is simple, affordable, and readily available.  1.2 THESIS OBJECTIVE AND SIGNIFICANCE 1.2.1 Objectives The importance of having a readily available and cheap material for arsenic removal has been an important field of research in the past few years. The main objective of this study is to asses the potential for using natural iron oxide minerals to remove arsenic (As) from water. The research will also answer the following questions: •  Do various Fe-Ox minerals adsorb As?  •  How does adsorption onto Fe-Ox depend on pH?  •  What are the As adsorption mechanisms?  •  What is the stability of the final As -FeOx product?  •  Which of the studied Fe-Ox minerals looks the most promising for As removal from water?  •  Can a simple household method be developed for the removal of As from water?  1.2.2 Significance and Contribution of the Work A simple method was developed to remove arsenic from water using natural Fe-Ox minerals. The method involves mixing natural iron mineral bearing soils (lateritic soils) with arsenic contaminated water for ten minutes and then filtering off the As-loaded solids (e.g. through a coffee filter). Arsenic levels in the treated water were found to be  4  below the drinking water standards. The treatment method is cheap and simple, making it suitable for house-hold use. For the obvious socio-economic reasons, large technologically-advanced industrialized plants are not feasible in many villages of developing countries. Instead, technologies are required with which small communities and villages can decontaminate their own drinking water at the lowest cost possible. Moreover, the developed method has the following advantages: •  Iron oxides are readily available and abundant in most parts of the world.  •  Goethite (laterite) is the most chemically stable iron oxide.  •  There is no need for pH adjustment prior to water treatment, i.e. it can be carried out at natural pH.  •  The final product has a high degree of stability.  •  The method is simple, fast, and effective (less than lppb of As left from initial 20ppm).  •  The method is suitable for villages or communities with either spread population (i.e. South America) or non centralized water system (i.e. Bangladesh)  •  Electroacoustic tests, adsorption tests, and leaching tests data were used together successfully to get a better understanding of the adsorption mechanisms of arsenic onto Fe-Ox.  5  2. LITERATURE REVIEW 2.1 ARSENIC IN THE ENVIRONMENT 2.1.1 Arsenic in Drinking Water . National and International Standards for Drinking Water Currently, the drinking water standard for arsenic is 50 ppb in most of countries. Table 2.1 presents selected arsenic regulatory limits. Table 2.1 Arsenic Regulatory Limits Organization Canadian Federal-Provincial-Territorial Committee on  Date  Limit  April, 2004  25ppb  January, 2006  lOppb  Drinking Water USA Environmental Protection Agency (USEPA) Argentina National Food Code  50 ppb World Distribution of Groundwater Arsenic Problems Concentrations of arsenic in fresh water vary depending on the source of As, the amount available and the local geochemical environment. A number of large aquifers, potential sources of drinking water, in various parts of the world have been identified as contaminated with arsenic (Fig.2.1).  6  FAIRBANKS.  a l e ij r./.N ISLANDS  #c:  •  SOUTH-WEST ENGLAND  NOHfHCENTRAL MEXICO  HALIFAX. OVA SCOTIA  KAMCHATKA  | . tAr*J*L KYUSHU. .M»AN  ::: PLE>  I AiillN R S I M6XSCX  1  .  •  7IMAPAN • : IA VIFf|iAf.S  EL S A L V A D O R #  I I >'. >  -  ASHANTI REGION. GHANA  SBABWE ANTOFAGASlAl CHILE N T  A"-*"  »-  A  SOUTH A F R I C A  LEGEND M  CHACO-PAMPEAN RAIN, ARGENTINA  WAIHAKEI, W W ZEALAND  Arsenic-affected aquifers  •  Arsenic rslatsd to mining operations  •  Geotrwrmai waters  Fig 2.1 Distribution of Documented World Problems with As in Groundwater (Smedley & Kinniburgh, 2002) Arsenic in B. C. Information regarding arsenic levels in Canadian drinking water is fragmentary. Currently, a research group from the University of Alberta is working on creating an arsenic map for water in Canada. In B.C., two of the most relevant studies are the survey performed by the Coast Garibaldi Health Unit, and the one carried out by the Ministry of Water, Land and Air Protection through the Water Quality Check Program. In 1994, the Coast Garibaldi Health Unit (BC) initiated a large scale well water survey, with the objective to asses the extent of arsenic concentrations in wells in the Powell River and Sunshine Coast areas, and to identify geographic locations with high arsenic levels (Carmichael, 1995). A total of 199 wells in Powel River and 259 wells in the Sunshine Coast area were sampled and analyzed for dissolved arsenic. In general, the deepest wells sampled were  7  the wells with the highest arsenic concentration. In the Powell River Region, 11% of the wells had arsenic in excess, while none of the wells sampled in the Powell River North area had arsenic values exceeding the guidelines. In the Sunshine Coast Region 21% of the sampled locations had arsenic in excess, and none of the wells in the Gibson Area had arsenic values exceeding the guidelines (Carmichael, 1995). Moreover, the Ministry of Water, Land and Air Protection evaluated the results of groundwater samples obtained through the Water Quality Check Program, carried out between 1993 and 1997. Over 2,100 samples were analyzed for arsenic, of which 2.2% had arsenic levels over 25ppb. Arsenic concentrations above this level were found in some rural wells near the communities of 100 Mile House, Bowen Island, Burns Lake, Chase, Kamloops, Quesnel, the Sunshine Coast, Vanderhoof, Vernon and Williams Lake. High arsenic levels have been found in isolated wells of Saltspring Island, the Lower Mainland, and near Nukko Lake. Arsenic levels above the drinking water guideline may also occur locally in other regions of the province (British Columbia Ministry of Health Services, 2002). In addition, there are some mines sites in BC with elevated arsenic in the wastes and drainage, some examples include: Equity Silver, Eskay Creek, Golden Bear, Mt Washington, Snip and Sullivan (Arsenic Technical Workshop Proceedings, 2002). Health effects associated with arsenic in drinkins water Humans are exposed to arsenic from soil, water, air, and food. Arsenic is rarely present in the free state (As ) in the environment, but is widely distributed as both, inorganic and 0  organic compounds. Arsenic exists in the -3, +3 and +5 oxidation states. Toxicological  8  evidence showed that, both forms of arsenic, +3 and -3, are more toxic than the pentavalent forms, and inorganic compounds are more toxic than organic compounds (ICPS, 2001). Effects of arsenic on human health can be acute, due to consumption of arsenic at high doses, or chronic, due to a long time exposure to low doses (Roy and Saha, 2002). People drinking water contaminated with arsenic are subject to chronic effects. Long term exposure to arsenic-contaminated water may lead to various diseases such as conjunctivitis, hyperkeratosis, hyperpigmentation, cardiovascular diseases, disturbance in the peripheral vascular and nervous system, skin cancer, gangrene, leucomelonisis (WHO, 1981). Effects on the lungs, uterus, genitourinary tract and other parts of the body have been detected in the advanced stages of exposure to arsenic. High concentrations of arsenic in drinking water also result in an increase in stillbirth and spontaneous abortions (Roy and Saha, 2002). Some examples of chronic effects due to drinking water contaminated with arsenic are: •  Arsenic contamination in ground water of Taiwan resulted in black-foot disease (Jain & Ali, 2000).  •  In Antofagasta, Chile, over 12% of the population exhibit dermatological manifestations related to arsenic due to consumption of high arsenic containing drinking water (Jain &Ali, 2000).  •  Endemic arsenism was discovered at the Tianshan Mountains in Xinjian, China. Some indices, such as the rate of keratosis on palm, pigmentation or colorless spot on patients, correlate well with the annual intake of arsenic and the cumulative arsenic exposure (mg/L of arsenic in years). The high arsenic intake also resulted in  9  increasing rates of malignant tumor, hypertension, and shortened life span (Wang Lianfang and Wan Shenling, 2003). •  Studies performed in Argentina, showed that drinking water with high arsenic content (higher than 100 ppb) promoted the development of bladder and prostate cancer (Lerda, 2003).  Fig. 2.2 presents different sources of human exposure to arsenic and various modes of arsenic toxicity. Drinking Water/ Foods / Drugs  \ y-=  1  Agricultural Mminij.Co il/Buriiing  Natural Dopo^its Soil  Human Exposure to Arsenic  Uptake/absorption Excretion Accumulation (skin, hair)  Cellular Metabolism Toxicity  Ncn-carcinogenic  Carcinogenic  Cardiovascular and Non-cardiovascular Effects Chromosomal Abnormalities  Modification Of Cell Proliferation  Aberrations In Gene Expression  Altered DNA Repair  Oxidative Stress  Fig. 2.2 Sources of Human Exposure to Arsenic and Various Modes ofArsenic Toxicity (Roy & Saha, 2002)  10  2.1.2. Sources ofArsenic in the Environment Natural Sources ofArsenic Arsenic occurs in the environment in rocks, soil, water, air, and in biota. Arsenic is the twentieth most abundant element in the earth's crust (USEPA, 2000). Concentrations of As in the earth's crust vary from 0.1 to several hundred ppm, depending of the types of rocks being considered (Cullen & Reimer, 1989). Concentrations of arsenic in a variety of environmental media are presented in Table 2.2. Table 2.2 Arsenic Concentrations in Environmental Media (Smedley&Kinniburgh, 2002) Environmental Media Arsenic Concentration Range Air  1.5-53 ng/m  Rain from terrestrial air  0.013-0.032 ug/L  Rivers  0.83 ug/L  Lakes (B.C.)  0.28 (O.2-0.42) ug/L  Ground water -Baseline UK  <0.5-10ug/L  Ground water -mining contaminated  50-10,000 ug/L  Sea water- Deep Pacific and Atlantic  1.0-1.8 ug/L  Soils  0.1-55mg/kg  Stream/river sediment  0.6-50 mg/kg  Lake sediment-BC  0.9-44 mg/kg  Igneous rock  0.2-113 mg/kg  Metamorphic rock  0.1-143 mg/kg  Sedimentary rock  3-500 mg/kg  3  Arsenic is a major constituent of many mineral species in igneous and sedimentary rocks. Among igneous rock types, the highest arsenic concentrations are found in basalts. Sedimentary rocks particularly iron and manganese ores often contain higher average arsenic concentrations than igneous rocks (USEPA, 2000).  11  Arsenopyrite, (FeAsS), realgar (AsS), and orpiment (AS2S3), are the most important minerals containing arsenic, and they are commonly present in sulfide ores of other metals including coopper, lead, silver and gold. Arsenic may be released from these ores to soil, surface water, ground water and the atmosphere. Table 2.3 lists common minerals that contain arsenic. Table 2.3 Common Minerals that Contain Arsenic (USEPA, 2000) Arsenopyrite, FeAsS Realgar, AS4S4 Smalite, C0AS2  Tennanntite, 4CU2SAS25S3  Lollingite, FeAs2  Chloanthite, N i A s 2  Cobaltite, CoAsS  Proustite, 3Ag2SAs2S3  Orpiment, AS2S3  Niciolite, NiAs  Gersdoffite, NiAsS  Enagirte, 3CU2SAS2S5  Concentration of arsenic in water is highly variable (Table 2.2). It depends on the source of arsenic, the amount available and the local geochemical environment. Under natural conditions, the greatest range and the highest concentrations of arsenic are found in groundwater as a result of the strong influence of water-rock interactions and the greater tendency in aquifer conditions (physical and geochemical) to be favorable for arsenic mobilization and accumulation (Smedley & Kinniburgh, 2002). In addition, geothermal waters can be sources of arsenic in surface and ground water (EPA, 2000). Arsenic concentrations in soils depend in part on the parent materials from which the soils were derived, although they may be enriched by other sources, such as anthropogenic sources. In addition, very high natural concentrations of arsenic can be found in soils that overlay deposits of sulphide ores (Smedley & Kinniburgh, 2002).  12  Arsenic content in sediments depends on texture and mineralogy. As the major arsenic binding mineral in sediments are metal oxides, (Fe, Al and Mn), elevated concentrations of arsenic tend to reflect the content of those oxides. In addition, high concentrations of arsenic are also common in mineralized areas (Smedley & Kinniburgh, 2002). Arsenic Anthropogenic Sources Men activities release arsenic into the air, water and soil, at the end, these emissions can affect residue levels in plants and animals. The major current and past anthropogenic sources of arsenic, which are wood preservatives, agricultural uses, industry, and mining and smelting are shown in Table 2.4. Table 2.4 Summary of Current and Past Uses ofArsenic (USEPA, 2000) Sector Uses Lumber  Wood preservatives  Agriculture  Pesticides, insecticides, defoliants, debarking agents, soil sterilant  Livestock  Feed additives, disease preventatives, animal dips, algaecides  Medicine  Antisyphilitic drugs, treatment of trypanosomiasis, amebiasis, sleeping sickness  Industry  Glassware, antifouling  electrophotography, paints,  dye  and  catalysts, soaps,  pyrotechnics,  ceramics,  mining,  pharmaceutical substances, alloys (automotive solder and radiators), battery plates, solar cells, optoelectronic devices, semiconductor applications, diodes in digital watches  Earlier, arsenic was widely used for preparation of pesticides, insecticides and food additives. These uses are currently declining (USEPA, 2000).  13  Arsenic and arsenic compounds are used in a variety of industrial applications. Most significant sources of arsenic in industry are listed in Table 2.4. Arsenic metal is used in the production of lead-acid batteries. In addition, primary and secondary lead smelters, primary cooper smelters, and secondary aluminum operations are potential sources of arsenic. Also, high concentrations of arsenic may occur in areas that are near or affected by current or historical mining activities (USEPA, 2000). 2.1.3 Environmental Transfer ofArsenic Arsenic Transport between Media Arsenic is emitted into the atmosphere by natural processes and by men activities. Natural processes include volcanic activity, erosion, sea spray, forest fires and lowtemperature volatilization. Some of the anthropogenic sources are related to smelting operations and fossil fuel combustion. Estimates have placed the ratio of natural to anthropogenic atmospheric inputs at 60:40, with natural emissions of arsenic largely comprising of low-temperature volatilization from soils (60%), and anthropogenic emissions dominated by metal production (Cullen and Reimer, 1989). Fig 2.3 shows the arsenic cycle in the environment. Arsenic is released into the atmosphere primarily as AS2O3, and exists mainly in the form of particulate matter. These particles are dispersed by the wind to a varying extent depending on their size, and are returned to the earth by wet or dry deposition. In addition, arsines that are released from microbial sources in soils or sediments undergo oxidation in the air, reconverting the arsenic to less volatile forms that settle back to the ground (WHO, 2005).  14  Fig 2.3 As Cycle. Environmental transfer ofArsenic (EPA, 2000) The dissolved forms of arsenic in water include arsenate, arsenite, monomethylarsonic acid (MMA) and dimethylarsinic acid (DMA). As(III) and As(V) species can change oxidation states depending on Eh, pH and biological processes. Transport and partitioning of arsenic in water depends on the chemical form of arsenic and on interactions with sediments. Some arsenic species have an affinity for certain mineral surfaces (Fe- oxides, clays, A l - oxides) and organic matter, and this can affect their environmental behavior. The distribution and transport of arsenic in sediment is a complex process that depends on water quality, native biota and sediment type (WHO, 2005). Arsenite dominates in soil under reducing conditions, while arsenate would be present in well-drained soils as  H2ASO4  if the soil was acidic, or as H A s 0 " if the soil was z  4  alkaline. Oxidation, reduction, adsorption, dissolution, precipitation and volatilization reactions of arsenic commonly occur in soil. Arsenic from weathered rock and soil may  15  be transported by wind or water erosion. Rainwater or snowmelt may leach soluble forms into surface water or groundwater. However, since many arsenic compounds tend to adsorb to soils, leaching usually results in transportation over only short distances in soil. In addition, soil microorganisms may reduce arsenic compounds to volatile forms (arsines) that then enter the atmosphere (WHO 2005). Moreover, three major modes of arsenic biotransformation have been found to occur in the environment: redox transformation between arsenite and arsenate, the reduction and methylation of arsenic, and the biosynthesis of organoarsenic compounds. There is biogeochemical cycling of compounds formed from these process (WHO 2005). Arsenic in Soils As Chemical Forms in Soils Total As concentration in a soil includes both organic and inorganic As. Usually, organoarsenical complexes constitute a minor fraction of the total As in soil. Inorganic arsenic species in soil are more significant. Their presence and concentration depend on both soil pH and redox conditions (pe, where pe is obtained by dividing Eh by 59.2) (Sadiq, 1995). At pH 7, As (V) species are more abundant in oxic soils conditions that are characterized by pe + pH >10, according to the following sequence: HAs0 " > H As0 " > A s 0 " > H As0 ° 2  3  4  2  4  4  3  4  (Sadiq, 1995)  On the other hand, As (Ill)-form at pH 7, are expected to be found in relatively anoxic soils with pEh + pH <7, according to the following sequence: HAs0 ° = H As0 ° >As0 ' = H As0 " 2  3  3  2  2  3  (Sadiq, 1995)  16  Figure 2.4 shows the chemical forms of arsenic and their transformation in soils. Arsenate  FeAs0  4  Reduction  Arsenite  As S 2  bacteria  I Dimethyl arsine (volatile, v.toxic)  3  I Trimetnyl arsine (volatile, v.toxic)  Fig 2.4 Chemicalforms ofArsenic and their transformation in soils (Bhumbla & Keefer, 1994) Arsenic Adsorption and Precipitation on Soils Adsorption on soil colloids is an important scavenging mechanism. The adsorption capacity and behavior of these colloids (clay, oxides or hydroxides of Al, Fe and Mn, calcium carbonates and /or organic matter) are dependent on factors such as hydration, soil pH, and crystallinity. In general, Fe oxides/hydroxides are most commonly involved in the adsorption of As in both acidic and alkaline soils. The surfaces of Al oxides/ hydroxides and clays may play a role in As adsorption, but only in acidic soils. The carbonate minerals are expected to adsorb As in calcareous soils. The role of Mn oxides and biogenic particles in the As adsorption process in soils appears to be limited to acidic soils. Kinetically, As adsorption may reach over 90% completion within hours (Sadiq, 1995).  17  Precipitation of a solid phase is another mechanism of As removal from soil. In acidic oxic and suboxic soils Fe-arsenate (Fe3 (AsO^) may control arsenic solubility, whereas in the anoxic soils, sulfides of As (III) may control the concentrations of the dissolved As in soil. In alkaline acidic oxic and suboxic soils, precipitation of both Fe and Ca arsenate may limit As concentrations in soil (Sadiq, 1995). Field observations suggest that direct precipitation of discrete As solid phase may not occur, except in contaminated soils. However, secondary precipitation of As compounds may occur on soil colloid surfaces subsequent to its adsorption. In the secondary precipitation, the first step would probably be the adsorption of dissolved As species onto soil colloidal surfaces. The adsorbed As ions would gradually and continuously move inside the hydrated layer of colloids. As a result of translocation, arsenic concentrations inside the hydrated mineral layer, with the passage of time, may rise to a level to precipitate as an As solid phase (Sadiq, 1995). Biological Transformation of Arsenic in Soils Arsenic volatilization is another scavenging mechanism operating in soils. Many organisms are capable of converting arsenate and arsenite, through biotransformation. Biotransformation of soils implies a series of redox reactions, converting arsenate and arsenite to several reduced forms, largely methylated arsines which are volatile; soils microorganisms involved in the biotransformation of arsenic in soils include fungi and bacteria (Sadiq, 1995). 2.1.3. 3 Arsenic in Water Arsenic can occur in the environment in several oxidation states but in natural waters is mostly found in inorganic form as trivalent arsenite (As(III)) or pentavalent arsenate  18  (As(V)). Organic As forms (monomethylarsonic acid, M M A A , and dimethylarsinic acid DMAA) may be produced by biological activity (biomethylation), mostly in surface waters, but are rarely quantitatively important (Smedley & Kinniburgh, 2002). Redox potential (Eh) and pH control the two main characteristics of arsenic in water: •  The high As concentration in groundwater as a consequence of its high solubility at natural water pH.  •  The high mobility of As, especially under reducing conditions.  The solubility of arsenic species in water depends on a number of factors, including pH, cations present, and adsorbing surfaces. Under certain conditions, the solubility and adsorption appear to be controlled by the presence of iron, manganese and their respective oxides (Driehaus et.al, 1995). At near-neutral pH, typical of most groundwaters, arsenate tends to become less strongly sorbed, thus it can be found in natural waters at relatively high concentrations (tens of u.g/L). Thus, arsenic is one of the most common contaminants in groundwater (Dzombak and Morel, 1990). Relative to the other oxyanion forming elements, arsenic is among the most problematic in the environment because of its relative mobility over a wide range of redox conditions. Arsenic is distinctive in being relatively mobile under reducing conditions. The oxidation of adsorbed As(III) to As (V) on mineral surfaces is an important factor in controlling the overall mobility of As. As (III) occurs as an uncharged molecule  ( H 3 A S O 3  0  )  in most  natural waters, which is less strongly adsorbed on most mineral surfaces than the negatively charged As(V) oxyanions (Manning and Goldberg, 1997). The distribution of species as a function of Eh-pH are shown in Fig.2.5  19  0  2  4  6  S  10  12  \^14  Fig2.5. Eh - pH diagram for the system As-O-H-S (Vink, 1995) Under oxidizing conditions the dominant form of arsenic (V) is H2ASO4" at low pH (less than about pH 6.9), while at higher pH, HASO4 " becomes dominant. The other forms, 2  H3ASO4  0  and As0 " may be present in extremely acidic and alkaline conditions, 3  4  respectively. Arsenic(III) is the dominant form under reducing conditions. At pH values lower than about 9.2, As(III) will exist as H3ASO3 , the uncharged arsenite species. 0  In the presence of extremely high concentrations of reduced S, concentrations of dissolved As-sulfide species can be significant. Reducing acidic conditions favor 20  precipitation of orpiment (AS2S3), realgar (AsS) or other sulfide minerals containing coprecipitated As (Cullen and Reimer, 1989). Therefore, high arsenic levels in waters are not expected to occur where there is a high concentration offreesulfide ions. Distribution of Arsenic Species in Surface and Ground Water The redox reaction of the As (III)/As(V) system can be described by the following equation:  H A s 0 + 2H + 2e -> H A s 0 + H 0 E = +0.206 V Equation (2.1) +  3  4  3  3  2  Where E is the potential needed for 99% oxidation of As(III) to As(V) at pH 7. This potential indicates easy oxidation of As(III), even by dissolved oxygen. However, the kinetics of oxidation by O2 is very slow and the conversion rate can be only a few percent per week. It can be concluded from these data that reduced arsenic is found in reduced, oxygen -free, ground water (Jekel, 1994). On the other hand, As(V) should strongly dominate over As(III) in oxygenated waters. And under slightly reducing conditions and/or lower pH, As (III) acid becomes stable. In fact, both forms of arsenic have been found in natural oxygenated or anoxic water, which demonstrates that both arsenic forms can co-exist under the two conditions, although one is more dominant than the other (Masscheleyn et al, 1991). In groundwater the ratio of As(III) to As(V) can vary greatly as a result of variations in the abundance  of redox-active solids, especially organic C, the activity of  microorganisms and the extent of convection and diffusion of O2 from the atmosphere. In strongly reducing aquifers, As (III) typically dominates (Smedley & Kinniburgh, 2002).  21  Furthermore, in lake and river waters, concentrations and relative proportions of As(V) and As(III) vary according to changes in input sources, redox conditions and biological activity, with As(V) being the dominant species. In addition, As(III) levels may be maintained in oxic surface waters by biological reduction of As(V), particularly during summer months (Andreae & Andreae, 1989).  2.1.3. 4 Arsenic in Sediments In river and lakes, arsenic is predominantly bound to sediments. Arsenic in sediments mainly comes from suspended solids that have acquired their arsenic burden from water. The mobilization of arsenic is therefore closely related to its interaction with sediments. Arsenic concentrations in suspended solids and sediments can be more than 2000 times higher than in river water, indicating that the suspended solids are a good scavenging agent and a sink for arsenic (Nriagu, 1995). Understanding how arsenic is released to water requires knowledge of not only the source of arsenic but also of the geochemically dominant process that affects its flowpath. Competing chemical reactions affect both the speciation and partitioning of arsenic between the aqueous phase and the solid mineral phase of the aquifer matrix. Desorption and remobilization of arsenic from sediments is controlled by pH, Eh, and arsenic concentration in the interstitial water of sediments, as well as by the presence and type of the adsorbing mineral/solids (Schreiber et.al, 2000). Several theories have been proposed to explain the subsurface mobilization of arsenic. These include i) the oxidation of As- containing sulfides, ii) the release of As(V) from reduction of iron oxides,, iii) the release of arsenic from clay minerals, and vi) the  22  exchange of adsorbed As(V) with fertilizer phosphates and/or other ions. These are not mutually excluding processes, and they can be either abiotic or biotic (Nicholas et. al, 2003). i -Oxidation of As containing sulfides in sediment. The most important sources of arsenic in many natural waters are sulfide minerals, whereby in such cases, sulfide oxidation seems to be the dominant process controlling arsenic release to groundwater. Sulfide minerals, particularly arsenopyrite (FeAsS), and others such as orpiment  AS2S3,  and realgar  AS2S2,  oxidize rapidly on exposure to the  atmosphere releasing the arsenic for partitioning between water and others minerals. Both microbially mediated redox reactions and abiotic process are involved (Plant et al, 2005). Oxidation of sulfide minerals can occur naturally or as result of mining activity. Mine dewatering, ore roasting, and the redistribution of tailings in ponds and heaps enhance sulfide oxidation (Plant et al, 2005). The oxidation of FeAsS can be described by the following reaction  4FeAsS (s) + 13 0 + 6 H 0 2  2  4Fe + 4S0 " +4 As0 " + 12H +3  2  4  3  4  +  Equation (2.2)  Which involves the release of acid, arsenic, and sulfate as acid mine drainage (AMD). r  Further acidity is released by the oxidation of the Fe  and precipitation of HFO or  schwertamannite. These minerals readsorb some of the released arsenic, reducing dissolved arsenic concentrations, and may eventually lead to the formation of FeAs0 .2H 0 (Plant et al, 2005). 4  2  23  Some examples of how sulfides minerals interact with arsenic species in water are discussed by Schreiber et al. (2000), Kim et al. (2000), Aiuppa et al. (2003), and Moore, (1988) ii- Arsenic and Fe- oxyhydroxides minerals in sediments The close association between arsenic and iron in minerals is frequently reflected by their strong correlation in soils and sediments. The consistent appearance of arsenic in the MnFe oxide fraction of sediments suggest that coprecipitation and/or preferential adsorption with these oxides may be involved in the control of dissolved arsenic concentrations in the overlying water column. Moreover, extraction procedures investigating the distribution of arsenic in sediments made by several researchers found that most of the arsenic in sediments was associated with the Fe oxide phases (Cullen and Reimer, 1989). The release of As from Fe-oxyhydroxides by the reducing groundwater is applicable for sediments, where rapid consumption of O2 by the microbial metabolism of organic matter produces a strongly reducing aquatic system. The sorbed arsenic is released and reduced by organic matter. This process, microbial mediated, releases As, HCO3" and Fe , into +2  solution according to the equation (Nickson et al 2000):  4FeOOH + C H 0 + 7 H C 0 -> 4 Fe + 8HCO3" + 6 H 0 +2  2  2  3  2  Equation (2.3)  Other examples of the associations of arsenic with Fe-Ox minerals in water are analyzed by Smedley et. al (2003), Fabian et. al (20003), Agget & O'Brien (1985).  24  iii- Arsenic and clay minerals in sediments Adsorption, desorption, oxidation and dissolution of As at clay sites are very important processes for the attenuation and mobilization of As. As(V) is preferentially adsorbed over As(III) on illitic and kaolinitic clays. Oxidation of As(III) to As(V) may occur on the clay surfaces. And desorption of As from the clays is higher for As(III) than As(V) (Lin & Puis, 2000). Iv- The role of phosphate and other anions in competitive adsorption Phosphates in the aquatic system can be introduced by either phosphate-rich fertilizers or organic matter decomposition. Phosphorus and arsenic are both Group V elements and, thus have similar chemical properties. Nevertheless, their biogeochemical behaviors differ in many respects in natural aquatic systems. Phosphorus is often a biolimiting nutrient in the ocean. Conversely, arsenic can be deleterious to marine organisms, and most often is a health risk to human. Both phosphorus and arsenic are released to aquatic environment through the weathering of rocks and by various human activities including mining, industrial and agricultural uses. Iron hydroxides are strong sorbents for phosphate and arsenate in the aquatic environment. Competitive adsorption between arsenic and phosphate can be responsible for releasing As into groundwater. The competitive adsorption between phosphate and arsenate onto ferrihydrite was investigated by Jain & Loeppert (2000). The presence of phosphate resulted in a significant reduction in arsenate and arsenite adsorption by ferrihydrite, with strong dependence on pH and phosphate concentration. The effect of phosphate on arsenate adsorption was greater at high pH than at low pH, whereas the opposite trend was observed for arsenite.  25  In addition to phosphate, others common dissolved substances in groundwater, such as bicarbonate and silicate may reduce the sorption capacity of arsenic onto some minerals. (Harvey el tal, 2002, Kim et al 2000, Appelo et al 2002).  2.1.3. 5 The role of bacteria in arsenic mobilization The role of microorganisms in the hydrologic mobility of arsenic in drinking water aquifers is very complex. Factored into such complexity are the competing chemical reactions that affect both the speciation and the partitioning of arsenic between the aqueous phase and the solid mineral phase of the aquifer matrix. At that scenery, microorganisms probably play an essential role in both the direct reduction and oxidation of the arsenic species as well as the iron minerals contained in these aquifers. Certain microbes use arsenic oxyanions for energy generation, either by oxidizing arsenite or by respiring arsenate. In aquifers, these microbial reactions may mobilize arsenic from the solid to the aqueous phase, resulting in contaminated drinking water. Arsenic-metabolizing bacteria have an impact on the speciation and mobilization of arsenic in water. On the other hand, there are other biotic processes that also play an important role in arsenic mobilization. Some examples are microbial Fe(III) reduction, and microbial oxidation of C, processes that can reduce Fe and release As. As(V) serves as a nutrient for certain anaerobes. The reaction is energetically favorable when coupled with the oxidation of organic matter. Such microbes are referred to as dissimilatory arsenate reducing prokaryotes (DARPs). They have been isolated from freshwater sediments, estuaries, soda lakes, hot springs, gold mines, the gastrointestinal tracts of animals and subsurface aquifer materials from Bangladesh. They can attack  26  As(V) adsorbed on solid phases like ferrihydrite and alumina, and reduce the As(V) contained in oxidized minerals such as scorodite (OremLand & Stolz, 2003). The microbiological oxidation of As(III) to As(V) can also impact the mobility and speciation of arsenic in the environment. More than 30 strains have been reported. They include both heterotrophic arsenite oxidizers (HAOs) and chemolithoautrophic arsenite oxidizers (CAOs). For example, some of them couple the oxidation of arsenite (electron donor) to the reduction of either oxygen or nitrate and use the energy derived tofixCO2 into organic cellular material and achieve growth (OremLand & Stolz, 2003). The process can be described by the reaction:  H As0 " + N0 " - HAs0 " + N0 " + H AG = - 56.5 kJ/mol Equation (2.4) 2  2  3  3  4  +  0  2  2.1.4 Arsenic Removal in Drinking Water Summary of Existing Technologies For the treatment of drinking water, technology selection depends on several factors, such as existing systems, the need to treat other contaminants, and the size of the treatment system. Precipitation/coprecipitation are the technologies most commonly used to remove arsenic from drinking water (USEPA, 2002,). They can also be used to remove contaminants other than arsenic such as hardness or suspended solids. In addition, the size of a drinking water treatment system may affect the choice of technology. Precipitation/coprecipitation processes tend to be more complex, requiring more unit operations and greater operational expertise and monitoring, while adsorption and ion exchange units are usually less complex and require less operator expertise and monitoring. Therefore, operators of smaller drinking water treatment systems are more  27  likely  to  select  adsorption or  ion exchange  to  treat  arsenic  instead of  precipitation/coprecipitation. Furthermore, permeable reactive barriers, elctrokinetics, phytoremediation and biological treatments are emerging technologies with few or no treatment facilities operating at full scale. Currently, treatments in use for the removal of arsenic from contaminated water are: Precipitation/coprecipitation, membrane filtration, adsorption, and ion exchange. All of these methods can lower arsenic concentrations to below 50ppb. A l l of them are affected by arsenic valence state (more effective in removing As(V)), pH, and presence of other compounds (USEPA 2002). Precipitation / coprecipitation These processes involve the use of chemicals to transform dissolved contaminants into an insoluble solid or form another insoluble solid onto which dissolved contaminants are adsorbed. The solid is then removed from the liquid phase by clarification or filtration. The presence of sulfate could decrease arsenic removal and affect method performance (USEPA, 2002). Membrane Filtration This method separates contaminants from water by passing it through a semi-permeable barrier or membrane. The membrane allows some constituents to pass, while blocking others. This technology can reduce arsenic concentrations to less than 0.050 mg/L. It produces large residual volume and tends to be more expensive than other arsenic treatment technologies. The presence of solids (suspended and dissolved), organic  28  compounds and colloids may cause membrane fouling, and affect method performance (USEPA, 2002). Adsorption Treatment This method concentrates solutes at the surface of a sorbent, thereby reducing their concentration in the bulk liquid phase. The adsorption medium is usually packed into a column. As contaminated water is passed through the column, contaminants are adsorbed. The presence of suspended solids, (organics, silica or mica) may cause fouling, and affect adsorption performance (USEPA, 2002).  2.1. 4.4 Ion Exchange Treatment This method exchanges ions held electrostatically on the surface of a solid with ions of similar charge from solution. The ion exchange medium is usually packed into a column. Arsenic contaminated water is passed through the column and the contaminant is removed. The presence of organics, suspended solids, calcium or iron can cause fouling of ion exchange resins. The presence of Fe(III) could cause arsenic to form complexes with the iron and then not be removed by ion exchange (USEPA, 2002).  2.2 IRON OXIDES 2.2.1 Natural Iron Oxide Occurrence and Description Introduction Iron oxides are common naturally-occurring compounds. Initially, the formation of iron oxides (Fe-Ox) predominantly involves aerobic weathering of magmatic rocks, in both  29  marine and terrestrial environments. A redistribution process may follow either by erosion or reductive dissolution in the form of Fe(II), and then oxidative re-precipitation (Cornell & Schwetmann, 2003). Most common oxides, hydroxides or oxides-hydroxides are listed in Tables 2.5 & 2.6 Table2.5 Most common Iron Oxides (Cornell & Schwetmann., 2003) Name  Formula  Hematite  oc Fe 0  Magnetite  Fe 0 (Fell, Fe III, 0)  Wustite  FeO  2  3  3  4  Table 2.6 Most common Iron Oxides -Hydroxides and Hydroxides (Cornell & Schwetmann, 2003) Name  Formula  Goethite  oc FeOOH  Lepidocrocite  Y FeOOH  Akaganeste  P FeOOH  Shwertmannite  F e 0 (0H) (S0 ) .nH 0  Ferrihydrite  Fe H0 . 4H 0  Bernalite  Fe(OH)  16  16  y  5  4  8  x  2  2  3  In the most common compounds, iron is in the trivalent state (ferric). Iron oxides consist of closely packed arrays of anions (usually hexagonal or cubic) in which the interstices are partially filled with divalent or trivalent iron, predominantly in octahedral, but in some cases in tetrahedral coordination.  30  The oxides-hydroxides can be dehydroxylated to their oxide counterparts. Other characteristics include low solubility and high energy of crystallization. Due to the latter, they form only minute crystals in the environment. Iron Oxides and hydroxides therefore have high specific surface areas (Cornell & Schwetmann, 2003). Iron Oxides in Rocks and Ores Iron with an average concentration of 51 g/kg is the fourth most abundant element by weight after O (466g/kg), Si (269g/kg) and Al (81 g/kg) in the earth's crust. Most rocks contain iron oxide minerals of varying nature and abundance. Table 2.7 presents dominant occurrences of iron oxides in geological formations (Cornell & Schwetmann, 2003/ Table 2.7 Dominant Occurrences of the Fe-Ox in Geological Formations (Cornell & Schwetmann, 2003). Mineral  Occurrence  Magnetite  Ubiquitous in rocks  Hematite  Red beds, banded iron formation, laterite crusts, hot brines  Goethite  Rocks: ubiquitous in small concentrations in consolidated and uncolosolidated rocks of any age but to a lesser extent in Paleozoic and older rocks Ores: minette, oolithic rocks, laterite crusts, Trummererze, bog ores  Moreover, magnetite, hematite and goethite are widespread over the world, description and natural occurrence are as follow:  31  Magnetite, Fe^CXi Magnetite is a black, ferromagnetic mineral, containing both Fe (II) and Fe (III), and has an inverse spinel structure (Nesse, 2000). Magnetite is readily formed to hematite or altered to goethite by weathering environment. It is a common accessory mineral in igneous and metamorphic rocks and can become concentrated in heavy mineral sands. Magnetite can be a major constituent of a rock through magmatic segregation and in high temperature hydrothermal veins. It represents reducing conditions relative to hematite (Cornell and Schwetmann, 2003). Magnetite is one of the most abundant and widespread of iron oxides. It is found in diverse geological environments and in some deposits in sufficient abundance to constitute an important iron ore. Around 1-2% of the world's iron is mined from intrusive magmatic segregation ores. Kiruna (Sweden) is a typical deposit. The ore consists of fine grained magnetite or, more rarely, hematite. Others examples of igneous iron deposits are Pea Ridge and Iron Mountain (Missouri, USA), Gransgerberg (Sweden), Larap (Philippine), El Laco (Chile), Cerro de Mercado (Durango, Mexico) (Gilbert & Park, 1986). Hydrothermal magnetite deposits are found at many locations around the Pacific basin, Most notably they are in Chile (El Romeral and El Tofo), Peru (Marcona and Acari), Central America, Australia, and Japan (Gilbert & Park, 1986). Hematite oc Fe^Oa Hematite is stable in the weathering environment and is commonly produced by weathering. Like goethite it is extremely stable and is often the end member of transformation of other iron oxides, though it may be altered to iron hydroxide minerals.  32  Hematite is also named specularite or oligiste. It is widespread in rocks and soils; its color is blood red if finely divided and black or grey if coarsely crystalline. It is found with other Fe-Ti-0 minerals in igneous and metamorphic rocks. In sedimentary rocks it forms as large banded iron formations and is the cement in red sandstones (Cornell and Schwetmann, 2003). Hematite forms under a wide range of conditions. It occurs in soils, either derived directly from de decomposing underlying rocks or as a result of alteration of pyrite, magnetite, goethite, siderite, or iron silicates. Hematite forms most commonly from magnetite through martitization. High grade hematite ore are usually found in soils where the silica has been removed by subsequent solution. It accounts for the red color of soils throughout the world, especially in tropical regions (Lateritization), and in many Triassic sedimentary deposits. Banded iron formations (BIF), metamorphosed sedimentary deposits, are large hematite ore bodies, which are thick sequences over vast areas of alternating thin beds of hematite and silica. 90% of iron consumed is extracted from Precambrian cherty banded iron formations (Gilbert and Park, 1986). The largest and most abundant ore deposits are around the Atlantic Ocean and the Indian Ocean. Nevertheless, there are a few BIF deposits around the Pacific Basin. Major districts are present in Brazil, Venezuela, the Lake Superior region (USA), the Labrador region (Canada), in South Africa, along the west coast of Africa (Gabon, Liberia and Mauritania), Russia, India, Manchuria and Western Australia (Gilbert and Park, 1986). Major iron deposits, predominantly hematite, are found around Lake Superior in Minnesota and Wisconsin, in the southern Appalachians; in Ukraine, China and India; in  33  Australia, especially in Western Australia; in Africa, especially Liberia; in Venezuela in the Orinoco region, and in Brazil in the States of Minas Gerais and Parana (Dana, 1997/ Hematite is found in association with limonite, siderite, or chamosite in oolitic ferruginous deposits, sedimentary deposits of variable composition, such as the Clinton ores of the southeastern USA, the Minette ores of the Alsace-Lorraine (France) and the deposits at Wabana, along the coast of Newfoundland (Gilbert and Park, 1986). Goethite oc FeOOH Goethite is one of the most thermodynamically stable iron oxides at ambient temperature. It is a common mineral formed under oxidizing conditions as a weathering product of iron-bearing minerals. In massive crystal aggregates goethite is dark brown or black, whereas the powder is yellow and is the responsible for the color of many rocks, soils and ochre deposits. Goethite forms the gossans over metalliferous veins and deposit (Cornell & Schwetmann, 2003). "Brown iron ore" (=limonite) is a principal component of laterites, which are formed in countries with warm, humid climates where rock decay has progressed without interruption for a long time. Also is the principal component of common rust everywhere iron and steel objects have been subject to corrosion. Goethite is a very common mineral nearly everywhere. Some localities are: Lake Superior deposits, in pegmatites Colorado, fine "pseudormorphs of "limonite" (goethite) after pyrite, Utah. As residual brown iron ores in the Appalachian region, mainly in Alaska, Georgia, Virginia, and Tennessee (USA). Cornwal (England); Siegen (Germany). Widespread in residual iron ore deposits in Ukraine; and in Russia; Minas Gerais (Brazil) (Gilbert and Park, 1986).  34 Iron Oxides in Soils A generalized overview of the occurrence of goethite, hematite in soils is given in Table 2.8 Table 2.8 Summary of the Occurrence of Iron Oxides in Various Soils (Cornell & Schwetmann, 2003) Mineral Major Soils Goethite  Aerobic and anaerobic soils of all regions  Hematite  Aerobic soils of subtropical, Mediterranean and humid to subhumid tropical regions (lateritic and plinthitic soils, red Mediterranean soils, oxisols, ultisols) Usually absent in soils of temperate and cool regions  -Goethite in Soils Due to its high thermodynamic stability, goethite is by far the most common Fe oxide in soils. For this reason, soils containing goethite as the sole Fe oxide occur around the globe and predominate in cool to temperate, humid climates. Furthermore, goethite occurs in association with every other common Fe oxide. In warmer regions it is commonly associated with hematite, whereas in cooler climates ferrhydrite and lepidocrocite are frequent partners. Where evenly distributed within the profile and not masked by black humic matter, goethite imparts a yellow - brown color to the soil profile. It can also be concentrated locally in mottles, concretions, ferricretes and other forms of secondary Fe oxide accumulation (Cornell & Schwetmann, 2003). - Hematite and its Association with Goethite in Soils Hematite, having similar thermodynamic stability to goethite, is the second most frequent Fe oxide in soils, but in contrast to goethite, is restricted to soils in warmer,  35  predominantly subtropical and tropical climates. Soils of these zones are often bright red because of the red color of hematite. Hematite very rarely occurs as the sole oxide in a soil, but it is often in close association with goethite (Cornell & Schwetmann, 2003). -Magnetite in soils Lithogenic magnetite is a common mineral in the coarse, heavy mineral fraction of soils. In contrast, pedogenic magnetite has been discovered only very recently.  2.2.2 Iron-Arsenic Compounds Interaction in The Environment Arsenic mobility in the aqueous environment is greatly influenced by arsenic adsorption onto iron oxides and the presence of sulfur. In a sulfur and /or iron-free environment, arsenic is predicted to be very mobile under almost all conditions. Only under very strongly reducing conditions, be it acidic or alkaline, can native arsenic occur and be stable. When sulfur is considered, the formation of realgar  (AS2S2),  and orpiment  (AS2S3)  should  be taken into account. Realgar is an indicator of strongly reducing conditions, be it acidic or alkaline. Orpiment occurs in less reducing conditions, mainly acidic but partly also alkaline. Furthermore, in regard to the most common arsenic mineral occurrence in nature, FeAsS (arsenopyrite), iron has to be incorporated in the diagram. Both, thermodynamic data and the experimental evidence indicate that two of the most stable Fe oxide phases are goethite and hematite (Cornell & Schewertmann, 2003). Thus, ionic arsenate and arsenite species are combined with hematite in a large diagram field. Under reducing alkaline conditions realgar (AS2 S2) coexists with magnetite (Fe304), whereas arsenopyrite (FeAsS) is stable only under rather extreme reducing and alkaline  36  conditions. Scorodite, iron arsenate mineral (FeAsC>4.2H20), appears to be stable under relatively strongly acidic conditions which are also moderately to strongly oxidizing. The Eh-pH diagram for the As-S-Fe-EfiO system showing the most common interaction of As species in the aqueous environment is presented in Fig 2.6  Fig 2.6 Eh -pHDiagram for As -S-Fe-  H 0 System (Vink, 1996) 2  2.2.3 The Solubility ofFe- Oxides In general, the solubility of Fe (III) oxides is low while Fe(II) oxides are sparingly soluble. This means that except at extreme pH values, these compounds give a very low level of total Fe in solution. In the pH range 4 - 1 0 , and in the absence of complexing or 37  reducing agents concentration of Fe total is<10" M , though iron oxides dissolve slowly over a wide pH range (Cornell & Schwertmann, 2003). The solubility diagrams indicate that for all the iron oxides under consideration, there is a region of minimum solubility around pH 7-8, i.e. around the point of zero charge (pzc). As iron oxides are amphoteric, they dissolve in acid media to form cationic hydroxy species, and in basic media to form anionic, hydroxy species. Hence, solubility rises as the pH moves away in either direction, from the pzc (Cornell & Schwertmann, 2003). The solubility products of iron oxides (at 25°C) suggest that the order of solubility is ferrhydrite > goethite > hematite, but it may also be hematite>goethite since iron (II) oxides are more soluble than iron (III) oxides (Cornell & Schwertmann, 2003). Fig 2.7 shows the solubility for goethite, ferryhidrite and "soil - Fe". Curves were obtained using the appropriate equilibrium constants (Cornell & Schwertmann, 2003).  \  log a  Fe * 3  Fe  \ \  Fe(OH)*'  Fe(OH)« V \ V  •10  Fe(OH)',  lerrihydrite  /  ^ v r  " soil-Fe " goethite  -15  /  v  \ _ ^  y^—' /  8  10  12  14  Fig.2.7 Solubilities of Goethite, Ferrihydrite and "Soil-Fe" as a Function of pH (Cornell & Schwertmann, 2003)  38  2.2.4 The Adsorption Process and Iron Oxides Surface Chemistry Adsorption is the basis of most surface-chemical processes. Adsorption can be defined as "the net accumulation of matter at the interface between a solid phase and an aqueous solution phase" (Sposito, 1998). Moreover, as a result of adsorption, a charge may arise from chemical reactions at the solid surface. The charge of these particles is dependent on the medium's pH. Natural particles develop surface charge from (Sposito, 1998): •  Structural substitutions and disorder (intrinsic surface charge),  •  Reactions they undergo with ionic species in aqueous solution (adsorption  reactions), these reactions involve: -adsorption reactions with proton or hydroxyls (from water molecules dissociation) -adsorption with other ions (ligands) in solution. The central ion of a mineral surface acts as Lewis acid and exchanges its structural OH against other ligands (ligands exchange) Under dry conditions, surface Fe atoms may be coordinatively unsaturated. Because they carry unoccupied atomic orbitals, surface Fe atoms are Lewis acids (a Lewis acid is an electron pair acceptor) and react with Lewis bases (electron pair donor). In aqueous systems, therefore, they coordinate with hydroxyl ions or water molecules which share their electron pairs with Fe. Upon adsorption, the water molecules usually dissociate resulting in a surface covered by hydroxyl groups coordinated to the underlying Fe atoms. Hydroxylation of iron oxides is a fast reaction requiring minutes, or at most hours for completion. It is followed by further adsorption of water molecules which hydrogenbond to the surface OH groups (Cornell and Schwertmann, 2003).  39  The surface hydroxyl groups are the functional groups of iron oxides i.e. they are chemically reactive entities at the surface of the solid in an aqueous environment. They posses a double pair of electrons together with a dissociable hydrogen atom which enables them to react with both acid and bases. Iron oxides are, therefore, amphoteric (Cornell & Schwertmann, 2003). The charge on the oxide surface is established by dissociation (ionization) of the surface hydroxyl groups, i.e. adsorption or desorption of protons depending on the pH of the solution, represented by the following dissociation reactions (Cornell & Schwertmann, 2003)  EFeOH <-> E FeOH + H +  +  2  EFeOH <->EFeO" + H  +  Equation (2.5) Equation (2.6)  The Electrical Double Layer The charge on the solid that arises as a result of the ionization of the surface hydroxyl group is balanced by a layer of counter ions (from the electrolyte) of opposite charge located in the aqueous phase; this ensures that the interfacial region is electrically neutral. The charged surface, together with the layer of counter ions in the solution phase make up what is termed the electrical double layer (EDL). The separation of charges in the EDL results in the development of a difference in potential between the surface and the solution and the surface charge density, (J (coulombs m" ) can be related to a potential 2  gradient, IJJ (Cornell and Schwertmann, 2003).  40  Point of Zero Charge As a consequence of adsorption a positively charged surface or a negatively charged surface can be formed. Because these charges are pH dependent, at some critical pH value the net surface charge will be zero. This pH value is referred to as the point of zero charge (pzc). "The pzc is the pH value at which the net total particle charge is equal to zero". In general, iron oxides have a pzc in the pH range 6-10 (Sposito, 1998). Meanwhile, positive, negative and neutral functional groups can coexist on the oxide surface. At pH < pzc the FeOH2 groups predominate over the FeO" groups, i.e. although +  the surface has a net positive charge some FeO" groups are still present. At the pzc, the number of FeOHj* groups equals the number of FeO" groups, and as the pH increases, the number of FeO" groups increases (Cornell and Schwertmann, 2003). Adsorption of Anions on Fe-Ox The adsorption process involves interactions of the adsorbate (arsenate) with the adsorbent (iron oxides with surface hydroxyl groups). The metal ion (Fe) acts as a Lewis acid and exchanges the OH group for other ligands (anions) to form a surface complex. Anions may adsorb on Fe oxides either specifically or non specifically. Specific adsorption involves replacement of the surface hydroxyl groups by the adsorbing ligand (anion), i.e.  EFeOH + L"  <-» EFeL + OH"  E(FeOH) + L" <-> EFe L + 2 OH" +  2  2  Equation (2.7) Equation (2.8)  41  It is also termed chemisorption, inner sphere adsorption, and in the case of ligands, ligand exchange (Cornell and Schwertmann, 2003). Adsorption of ligands (such as  ASO4H2")  on the soil particles surface can take place via  three mechanisms (Stumm W., 1992), as is shown in Fig. 2.8 •  inner sphere surface complex  •  outer sphere surface complex, (both of them involve surface complexation  reactions) •  Adsorption in the diffuse-ion swarm, this adsorption mechanism neutralizes  surface charge only in a delocalized sense.  Fig 2.8 Surface Complex Formation of an Ion (e.g. Cation) on the Hydrous Oxide Surface. (Sposito, 1989)  42  Moreover, the diffuse-ion swarm and the outersphere surface complex mechanisms of adsorption involve almost exclusively electrostatic bonding mechanisms, and therefore the corresponding adsorption products are less stable than inner sphere surface complexes, which have a large proportion of covalent character. In addition, the inner and outer sphere complexes differ from each other depending whether a water molecule is interposed or not. -inner sphere complex: if no water molecule is interposed between the functional group and the ion or molecule it binds, this is called specific adsorption -outer sphere complex: If at least one water molecule is interposed between the functional group and the bound ion or molecule. During specific adsorption, the anion is bound directly to the particle surface, and because of the binding of the anions the net surface charge decreases or becomes more negative. The pH where the fixed surface charge becomes zero, and the point of zero charge, pzc, is shifted to lower pH values. Hence, specifically adsorbed anions decrease the pH of the point of zero charge (Stumm, 1992). Specific adsorption may take place on a neutral surface, or even on one with the same charge as the adsorbing species. Specifically adsorbing ions modify the surface charge on the oxide, they are usually tightly bound and not easily displaced (Cornell and Schwertmann, 2003/ Non-specific adsorption is also termed outer sphere adsorption. It is dominated by electrostatic forces, and hence is influenced by the ionic strength of the system. The adsorbing species retains its primary hydratation shell, i.e. at least one water molecule is interposed between the anion and the surface (Cornell and Schwertmann, 2003/  43  For non-specifically adsorbing anions, adsorption is negligible above the pzc. When electrostatics govern adsorption (i.e. non specific adsorption), the surface must have an overall positive charge in order for anion adsorption to take place, hence the region of maximum adsorption lies for pH <pzc of the solid. In contrast, where specific adsorption is involved, an overall positive surface charge is not required (only FeOH2 and FeOH +  groups) which explains why adsorption can occur at pH above the pzc (Cornell and Schwertmann, 2003).  2.2.4. 2 Arsenic (V) Adsorption onto Iron Oxides Sorption of arsenic (V) anions on hydrous oxides can be described via ligand exchange reactions in which hydroxyl surface groups are replaced by the sorbing ions. Arsenate sorption onto iron oxide can be described by reactions such as (Dzombak and Morel, 1990):  EFeOH + As0 " + 3 H o EFeH As0 ° + H 0  Equation (2.11)  EFeOH + As0 " + 2 H f> EFeHAs0 " + H 0  Equation (2.12)  EFeOH + As0 " + H <-» EFeHAs0 ' + H 0  Equation (2.13)  EFeOH + As0 "  Equation (2.14)  0  3  +  4  0  2  3  4  3  4  0  +  2  2  4  3  4  2  +  4  0  4  2  <-» EFeOHAs0 ~ 3  4  Adsorption is one of the processes that controls the mobility and bioavailability of As. Studies that quantified and modeled As(V) and As(III) sorption onto amorphous iron oxides, goethite, and hematite have been performed. Results showed that both As (V) and As (III) sorb onto iron oxides. In addition, arsenic sorption behavior is dependent on  44  its oxidation state, pH, and the relative affinity, of the minerals for the arsenic species. The examples presented below are an attempt to elucidate the mechanisms of arsenic adsorption onto iron oxides. As(V) and As(III) sorption onto amorphous iron oxide(HFO), goethite and magnetite were compared in order to study arsenic mobility in the environment. Experimental data were modeled with a diffuse double layer surface complexation model. Sorption of As(V) onto HFO and goethite is more favorable than that of As(III) below pH 5-6, whereas above pH 7-8, As(III) has a higher affinity for the solids. In addition, phosphate competes for sorption sites. As(V) and As(III) sorption isotherms indicate similar surface site densities on the three oxides (Dixit & Hering,2003). The kinetics of arsenate and chromate adsorption/ desorption on goethite using a pressure-jump relaxation technique were studied by Grossl, et.al.(1997). Information provided by this technique was used to elucidate the fate of arsenate and chromate in natural environments. The proposed mechanism for the adsorption of arsenate and chromate on goethite is a two step process resulting in the formation of an inner-sphere bidentate surface complex. Arsenic adsorption on amorphous iron oxides as a function of solution pH, solution ionic strength, and redox state was investigated by Goldberg & Johnston(2000). In situ Raman and Fourier transform infrared (FTIR) spectroscopy methods were combined with sorption techniques, electrophoretic mobility measurements, and surface complexation modeling to study the interaction of As(III) and As (V) with amorphous oxide surfaces. The following mechanisms of As sorption to these surfaces were proposed: arsenate  45  forms an inner sphere complex on Fe oxide, while arsenite forms both inner and outer sphere surface complexes on amorphous Fe-oxide. Lenoble et.al. (2002) studied arsenic adsorption on simple materials such as goethite and amorphous iron hydroxide, and more complex matrices such as clay sustained with titanium (IV), iron (III) and aluminium (III). These sorbents were characterized by various methods: XRD, FTIR, BET, DTA/TGA, surface acidity and zetametry. Elimination of arsenite and arsenate as a function of pH was also studied. Arsenate elimination was favored at acidic pH, whereas optimal arsenite elimination was obtained at 4<pH<9. Klaus et. al. (1998) compared the adsorption behavior of arsenite and arsenate on ferrihydrite, with regard to adsorption kinetics, adsorption isotherms, and the influence of pH on adsorption. Arsenate adsorption was higher at pH 4.6 than at pH 9.2, while As (III) adsorption was less influenced by the pH. Moreover, adsorption envelopes (graph of anion adsorbed at the same concentration vs. pH) exhibited broad adsorption maxima extending from approximately pH 6.8 to pH 9.4 for arsenite and from pH 5.2 to pH 7.0 for arsenate. The adsorption of arsenic onto hydrous ferric oxide (HFO) was examined by Wilkie & Hering (1996), for both As (III) and As (V). The effects of adsorbate / adsorbent ratios and of the presence of sulfate and calcium as co-occurring solutes were also examined. The observed results were compared with the results obtained using a surface complexation model to elucidate the factors affecting arsenic adsorption. For As (III), increased adsorption was observed with decreasing total arsenic / total iron ratios. Significant.competitive effects of sulfate on adsorption of As (III) and, to lesser extent for  46  As (V) were observed. The extent of competition was found to be pH-dependent. Cooperative effects of co-occurring solutes were also demonstrated. For example, As (V) adsorption at high pH was enhanced by calcium. This enhancement could be attributed to favorable electrostatics effects at the oxide surface.  '  In addition, the adsorption and mobility of As (III) and As (V) on an oxisol and its mineral constituents were studied as a part of a broader project aimed at selecting a soil liner to be used in tailing dams at a sulfuric gold ore plant (Ladeira and Ciminelli, 2003). Among the soil constituents, goethite was the most efficient adsorbent with regard to arsenic adsorption. Results showed that As (III) is more weakly bound than As (V), thus is more mobile, and that the oxisol and goethite were superior to gibbsite with respect to As immobilization. Adsorption and mobility were also correlated with electrophoretic mobility. Results indicated that As (V) is mainly adsorbed as an inner sphere complex, and As (III) may be adsorbed as either an inner or outer sphere complex. Iron oxides have also been investigated for their use as adsorbents for arsenic removal in water, for example synthetic goethite (Matis et.al, 1999), iron oxide-coated polymeric material (Katsoyiannis & Zouboulis, 2002), granulated iron hydroxide (GEH) (Wingrich & Wolf, 2002). But few studies have been carried out on the feasibility of the use of naturally-occurring iron oxides as adsorbents for arsenic removal; they are more attractive than the synthetic oxides because they are more cost effective. Simeonova (2000) used natural hematite in a pilot test in Mexico and a residual concentration below 0.05mg/L was achieved. Zhang W. et. al. (2004) focused their work on studying how various experimental parameters influence arsenic adsorption onto natural iron ores. Iron ore containing mostly hematite was found to be very effective for arsenic adsorption. As  47  (V) was lowered from 1 mg/L to below 0.01 mg/L in the optimum pH range 4.5 - 6.5 by using a 5g/L adsorbent dose. The adsorption capacity was estimated to be 0.4 mg As (V)/g adsorbent.  2.2.5 Conclusions Arsenic groundwater contamination is a serious health-related-problem all over the world; it affects millions of people due to the toxic effects of arsenic compounds. Currently known conventional technologies for arsenic removal in water, are not suitable for small villages of developing countries. Thus, there is a clear need for a simple, affordable, treatment that does not require water pre-treatment (i.e. pH adjustment), this treatment has to be able to meet water requirements for drinking and food preparation, ( 2 - 3 liters per capita per day) at low arsenic concentrations. On the other hand, the natural affinity of Fe-Ox for arsenic makes them a very suitable means for arsenic removal in water. Iron oxides are abundant in nature, cheap, appear to be an appropriate material to be used in an effective, reliable treatment for arsenic removal.  48  3. EXPERIMENTAL PROCEDURES  3.1 MATERIALS 3.1.1 Solids The iron oxide samples used in this study were goethite (Yellow Iron Oxide Pigment, YP2100) and magnetite (Black Iron Oxide, BK5099) from Elements Pigment Inc, Toronto, Ontario. Hematite was obtained from Lead South Dakota (David New Minerals, Hamilton Montana), and Laterite 1 and Laterite 2 were natural laterite ore samples from Highland Pacific's Ramu nickel laterite deposit in Papua New Guinea. Fe-Ox slurries were prepared by mixing an appropriate weight of Fe-Ox with distilled water.  3.1.2 Solutions Stock solutions containing 5 g/L of As (V) were prepared on a daily basis from sodium arsenate dibasic heptahydrate: (Na2HAs04.7H20) (Sigma ACS reagent) and distilled water. Experimental solutions were prepared freshly from these stock solutions as required. Dilute nitric acid (1M) and sodium hydroxide (2M) solutions were used as pH modifiers. They were prepared from a concentrated nitric acid solution (69.5%, Fluka, trace select) and solid sodium hydroxide (Fisher Scientific). Buffer pH solutions, 4, 7, and 10, were used to calibrate a portable Orion pH-meter. A leaching solution of magnesium chloride was prepared from magnesium chloride hexahydrate (MgCl .6 H 0 , Fisher ACS reagent). 2  2  49  Final arsenic solutions, from scoping, adsorption, leaching and water treatment tests were filtered through a 0.45 um pore membrane filter, acidified with FINO3, (pH<3), and sent to an analytical lab for arsenic analysis. As was analyzed by Enhanced ICP-MS (MLD 0.5ppb) for dissolved As in water samples.  3.2 PROCEDURES 3.2.1 Preliminary Tests Sample Preparation Laterite, magnetite, and goethite powders were used as received. Hematite wasfirsthandground, and then mortar-ground until a convenient fraction for electroacoustic analysis (< 1 mm) was obtained. The particle size distribution of each sample was determined with the use of a Malvern Mastersizer 2000. The specific surface areas of the powders were measured using a Quantasorb (Quantachrome) Bruauer, Emmet, and Teller (BET) analyzer. Results from particle size analysis and surface area are shown at Table 3.1. Mineralogical compositions of the samples, determined by XRD, are shown in Table 3.2. Table 3.1 Solids Particle Size & Surface Area Analysis Solid Specific Surface Area Particle Size analysis  -10 urn  m /g 2  (um)  %  Magnetite  6.58  89% < 11.7 pm  83  Hematite  3.77  89%< 74 um  33  Goethite  12.1  89% < 9.7 um  90  Laterite (#2)  81.2  89% < 56 um  57  50  Table 3.2 Mineralogical Composition of Samples {Elements Pigment Inc, Toronto, David New Minerals. Hamilton, Klein & Hallbom, 2002) Test Sample Mineralogical Composition Magnetite  98% Magnetite, (minor Hematite)  Hematite  96 % to 98 % Hematite, ( minor Quartz)  Goethite  99% goethite  Laterite (#2)  91% Goethite, 2.2% Talc, 1.7% Quartz, 4.9 % Gibbsite  Laterite (#1)  82.5% Goethite, 12% Talc, 4.5% Gibbsite  3.2.2 Scoping Tests A preliminary adsorption test for each sample was performed to estimate the suitability of different iron oxides for arsenic removal. The tests were carried out on the Fe-Ox samples. All the samples were tested under the same conditions of pH, temperature, arsenic concentration and time. The experimental parameters were chosen from literature (Matis etal. 1998, Wilkie & Hering 1995). Slurry concentration for magnetite, goethite and laterites was lOg/lOOmL, while hematite had a concentration of 0.5g/100mL. The initial arsenic (V) concentration was 4.95mg/L. The samples were placed on a shaking table for 3h. The pH of the samples were kept constant at 5.0 and monitored by a portable pH-meter. After adsorption, filtration was performed to separate the solids from solution. Then, samples were acidified and refrigerated prior to analysis.  51  One standard sample and one blank sample prepared with H N O 3 were sent also for analysis. The results received from the lab were corrected for any volume change due to washing in the filtration step. 3.2.3 Electroacoustic Tests In order to investigate the effect of arsenic adsorption on the surface charge of the tested iron oxides a series of systematic electroacoustic tests were performed. A Zeta Probe (Colloidal Dynamics, Warwick, Rl) was used as the measuring instrument. The probe was cleaned, calibrated for pH and the mixer set at 150 rpm. The results were recorded as the raw ESA signal in mPa*m/V vs. pH. Background corrections for the contribution of the background electrolyte to the overall ESA signal were made on each test. Electroacoustic methods use the Electrokinetic Sonic Amplitude (ESA) effect for measuring the zeta potential and surface charge of fine particles (Cannon, 1993). In ESA measurements an alternating electric field, of known amplitude and frequency is applied to a suspension of fine, charged particles. As the particles oscillate in the applied electric field, the dynamic mobility/zeta potential of the particles is determined by the magnitude of particle motion. The velocity of this oscillatory motion is proportional to the surface charge of the particle. The raw ESA signal measured is given by (Cannon, 1993):  ESA=P/E (mPa*m/V)  Equation 3.1  P is the measured magnitude of the pressure wave (in mPascals) E is the applied electric field strength (V/m)  52  There are two main limitations to the electrokinetic measurements, one is the concentration of the background electrolyte, and the other is particle size of the solid phase (Fe-Ox) (Cannon, 1993). The main problem associated with electrolyte solution is related to the fact that electrolyte ions also generate their own ESA signals during electroacoustic measurements. Therefore, the raw ESA signal of a suspension consists of two contributions, that of the suspended particles and that of the suspending medium (background electrolyte solution). To obtain the zeta potential or surface charge of the particles, the signal from the background electrolyte must be subtracted from the raw ESA signal. The Zeta Probe operating software provides this option and allows the user to reanalyze the sample by subtracting the background signal. Samples tested were reanalyzed and background signals were subtracted in all the cases. Changes in the ESA signal due to electrolyte concentration turned out to be negligible. Limitations to particle size are related to the fact that larger particles may not produce a measurable electroacoustic signal. The zeta potential of the particles is determined by the magnitude of particle oscillatory motion. Oscillation of very small, sub-micron particles easily keep up with the frequency of the field, but the motion of large particles lags significantly behind the applied field and the phases of the two sinusoidal signals do not match. While the magnitude of the ESA signal decreases with particle size, the phase lag increases. In zeta potential measurements with the use of electroacoustic instruments, this mobility spectrum is obtained by measuring the ESA signal of a suspension over a range of applied field frequencies. The Zeta Probe works at seven frequencies in the range from 0.3 to 3MHz. Particles at the upper sizing limit must therefore move with a sufficient  53  velocity at the lowest frequency to produce a measurable electroacoustic signal. The frequency must also be low enough so that the phase lag of the largest particle is smaller than the limit of 45"(Cannon, 1993). As reported by the manufacturer, the operating frequency range of the Zeta Probe allows the instrument to determine particles sizes up to 10pm. For coarser materials the instrument must be "told" the size distribution by imputing d50 and d85 (assuming a lognormal distribution) to obtain zeta potential, provided that such coarse particles give a measurable signal (Pawlik, 2005). Since the tested systems contained coarse particles, it was decided to report only the raw ESA value rather than the z potential. It should be noted that the zeta potential is zero when ESA is zero, so the iep and the pzc do not change. Three different sets of electrokinetic tests were carried out: A. Hematite slurry titrated, at constant pH, with arsenic (V) solution of known concentration (Table 3.3). B. Iron oxides slurries, alone titrated with acid/base over the whole pH range (back and forward titration) (Table 3.4). C. Iron oxides slurries, loaded with an arsenic (V) titrated with acid/base over the whole pH range (back and forward titration) (Table 3.5). Tests conditions are detailed in Tables 3.3,4 and 5.  54  Table 3.3 Test Conditions for Hematite, Titrated at Constant p, with As (V) Solution Mineral  Slurry  As (V) as  concentrationw/w)  titrant  pH = Constant.  Figure #  (mg/L) Hematite  25g/250g  5.0  4.2  4.4  Hematite  25g/250g  5.0  9.8  4.4  Table 3.4 Test Conditions for Titration of Fe-Ox Slurries Alone Mineral  Slurry  As (V) added  Concentration( w/w)  (mg/L)  Hematite  25g/250g  None  4.2  Goethite  12.5g/250 g  None  4.5  Magnetite  12.5g/250g  None  4.1  Laterite #1  12.5g/250g  None  4.6  Laterite #2  12.5g/250g  None  4.7  Figure #  Table 3.5 Tests Conditions for Titration of Fe-Ox Slurries loaded with As (V) Mineral Hematite  Slurry  As (V) added  Concentration w/w)  (mg/L)  25g/250g  •  20.0  •  50.0  Figure # 4.2  Goethite  12.5g/250g  •  50.0  4.5  Magnetite  12.5g/250g  •  50.0  4.1  Laterite #1  12.5g/250g  •  50.0  4.6  Laterite #2  12.5g/250g  4.7  •  50.0  tt  •  100.0  u  •  200.0  •  250.0  a  55  3.2.4 Adsorption & Leaching Tests Leaching tests were performed in order to assess adsorption mechanisms, differentiate between weakly and strongly adsorbed arsenic, and evaluate the chemical stability of the adsorption product. A sequential extraction procedure, with the use of a 1M  MGCL2  solution designed  specifically for arsenic was closely followed to study the leaching of arsenic from hematite, goethite, laterite, and magnetite (Keon et.al, 2001). Sequential chemical extraction quantifies arsenic associated with the solid phase based on the chemical properties of the target binding phases; it requires little specialized instrumentation as compared with other methods (Keon et.al, 2001). The sequences of extraction steps in the sequential extraction procedure are specifically designated to divide As based on the nature of its associations with the solid phase. Moreover, because the method allows the quantification of easily mobilized, adsorbed arsenic, the application of this extraction procedure will yield practical information in assessing the potential mobility of As in sediments. For example, arsenic removed by the MgCl extractant could potentially be mobilized by an increase in the ionic strength of 2  the aqueous phase, as in the case of seawater intrusion (Keon et.al, 2001). From this method any arsenic leached by  MGCL2,  was treated as weakly (ionically/  electrostatically) bound arsenic (Keon et.al, 2001). Table 3.6 summarizes the sequential extraction conditions.  56  Table 3.6 Sequential Extraction Conditions {Keon et.al, 2001) Extractant  Target phase  1M MgCl , pH 8, 25 C U  2  Ionically bound As  Possible mechanism Anion exchange of CI for  two repetitions + one water  As  wash  complex formed  possible Mg -  As  Prior to extraction, arsenic adsorption tests at pH 5.2 were performed on all the solid samples. A 5g/L As (V) solution was added to solid Fe-Ox (2 g) to give a final slurry volume of 40 mL and an arsenic concentration of 50mg/L. The slurry concentration (i.e. 12.5 g solid/ 250 g solution, ratio lg solid: 20g solution), and arsenic concentration (50 mg/L) were kept the same as in the electroacoustic tests. After adsorption, the slurries were split into two sub-samples A & B (volume 20 mL each). •  On samples #A, no leaching with MgCL; was carried out. These samples were  diluted up to 50 mL with distilled water (d 1:2.5), centrifuged, filtered, acidified and sent to the analytical lab for dissolved arsenic analysis. •  Samples #B, were further split into two equal-volume sub-samples (B-l & B-2),  with each of the sub-samples subjected in parallel to the same sequential extraction procedure. In every leaching step, an excess amount of M g C l 2 was used to ensure that the reagent did not become exhausted (Keon et.al, 2001). Sub-samples B - l and B-2 were placed separately into two 50 mL polypropylene centrifuge tubes, a suitable weight of MgCb was added, pH brought to 8, and final volume made up to 50 mL. These suspensions were placed in a shaker for 2 hours, and  57  then centrifuged for 25 min. The supernatant was decanted using a Teflon type syringe (with care not to remove sediment) and filtered through a 0.45 urn pore filter, acidified and sent to analytical lab for dissolved arsenic analysis. The residual sediment of each sub-sample was further treated as follows (sub samples B#2, and B#3): MgCh was added, pH brought to 8 and final volume made up to 50 mL. Suspensions were shaken, centrifuged, decanted, filtered, acidified and sent for arsenic analysis in the same way twice more. The last step comprised of washing the sediment with water (subsamples B#4). Arsenic was analyzed by Enhanced ICP-MS. Reagent blank for each Fe-Ox were analyzed in parallel. Figure 3.1 summarizes the sequential extraction procedure.  58  Fig 3.1: Sequential Extraction Procedure Adsorption Tests Fe-Ox 2g solid :40mL solution As(V) 50 ppm  To As analysis A d 1:2.5  To extraction procedure B Extraction Steps Two vessels 0.5/10mL + MgCl , pH=8 fill up to 50mL 2 h shaking 2  Liquid, to As anal. centrifugation  -lOOmL Solid To next extraction step  Samples #B2,B3,B 4  59  3.2.5 Adsorption Isotherm Tests Adsorption tests were performed on laterite #2 at constant pH in order to obtain the adsorption isotherms, find a correlation between electrokinetic data with the adsorption results, and to asses the influence of pH on adsorption. The experimental procedure for the adsorption tests follows recommendations from the literature (Pierce & Moore (1980), Manning & Goldberg (1997), Dixit & Hering, (2003)). The adsorption tests were carried out by adding a constant volume of freshly prepared arsenic (V) 5 g/L solution to a variable weight of the laterites. The pH of solutions was maintained constant and monitored during the first four hours at pH 5.2, and pH 9.2. After constant mixing overnight at room temperature, samples were centrifuged, filtered, acidified and sent to the analytical lab for arsenic analysis Tables 3.7 (pH = 5.2) and 3.8 (pH = 9.2) show the adsorption test conditions. Table 3.7 Adsorption Isotherm Tests at pH =5.2 Sample # w (g of solid)  mL ofAs sol.  Final vol.  As initial  mgAs/g sol.  (5.0 g/L) added mL  mg/L  Initial  4.2  5.1  2.0  100.0  100.0  1.96  4.3  10.20  2.0  100.0  100.0  0.98  4.4  10.30  2.5  100.0  125.0  1.21  4.5  2.5  2.0  100.0  100.0  4.00  4.6  1.0  2.0  100.0.  100.0  10.0  60  Table 3.8 Adsorption Isotherm Tests at pH = 9.2 Sample # w (g of solid)  mL ofAs sol.  Final vol.  As initial  mgAs/g sol.  mL  mg/L  initial  (5.0 g/L) added 9.2  5  2.0  100.6  99.40  1.99  9.3  10.30  2.0  98.3  101.73  0.99  9.4  11  2.75  110  125.00  1.14  9.5  2.5  2.0  100  100.00  4.00  9.6  1.0  2.0  100  100.00  10.0  The samples were filtered (0.45 um), acidified (to pH<3,  HNO3),  and refrigerated prior  analysis. One standard sample and one replicate were also sent for analysis. Adsorption isotherms, as the relation between the equilibrium concentration of arsenate in solution and the quantity of arsenate adsorbed onto the Fe-Ox surface, were plotted in order to better describe the adsorption process. In addition, the Langmuir equation was used in order to calculate the maximum adsorption capacity of the laterite. The Langmuir adsorption equation can be expressed as (Sparks, 1995)  q= kCbl{\+kC)  Equation 3.2  Where q, is the amount of adsorption (adsorbate per unit mass of adsorbent), and C is the equilibrium or final adsorbate concentration, k is a constant related to the binding strength, and b is the maximum amount of adsorbate that can be adsorbed (monolayer coverage) (Sparks, 1995). After rearranging to a linear form, Equation 3.1 becomes C/q =l/kb + C/b  Equation 3.3  61  Plotting C/q vs. C, the slope is lib and the intercept llkb. From this, it is possible to calculate maximum sorption (b) values for soils (Sparks, 1995).  3.2.6 Evaluation of a Simple Water Treatment Process The test was performed in order to evaluate a household method for arsenic removal in water at natural pH. The experiment consisted of mixing (hand mixing) 100 mL of water containing 20 mg/L of arsenic with 5 g of Laterite #2 at pH 7. After mixing, the samples were filtered through a coffee filter paper (double), acidified and sent for As analysis. One of the samples, sample #T4, still contained residual laterite particles which were allowed to settle over a period of 4 hrs. Then, the supernatant was removed using a Teflon type syringe, acidified, and sent to lab for As analysis. Arsenic was analyzed by Enhanced ICP-MS. The test conditions are shown in Table 3.9 Table 3.9 Test Conditions for the Evaluation of a Simple Water Treatment Process Time  Sample#  15 sec  Tl  60 sec  T2  lOmin  T3  lh  T4  62  C  4. RESULTS AND DISCUSSION  4.1 SCOPING TESTS Table 4.1 presents the results from scoping tests to asses As (V) adsorption by different iron oxides at pH 5.0. The initial arsenic concentration was 4.95 mg/L.  Table 4.1 Scoping Tests: As Adsorbedfrom Different Fe-Ox Slurries at pH 5.0 Slurry  Magnetite (10gsolid/100ml_ Hematite (0.5 gsolid/100ml_ Goethite (10gsolid/100mL Laterite (#1& #2) (10gso!id/100ml_  As adsorbed  As adsorbed  As adsorbed  mg 0.495 (all the arsenic)  mg/g solid  mg/m solid  0.0495  0.0075  0.103 0.495 (all the arsenic) 0.495 (all the arsenic)  0.206  0.055  0.0495  0.0041  0.0495  0.00061  2  The results from the scoping test presented in Table 4.1 show that all the Fe-Ox studied i.e., hematite, magnetite, goethite and laterite, are suitable as arsenic adsorbents. High arsenic uptakes per gram of solid can be achieved using magnetite, goethite and laterite. Hematite exhibited the lowest arsenic uptake per gram of solid. Most of Fe-ox, except hematite, did not reach saturation point due to their high solids concentration and low arsenic load. Hematite was the only one that did not adsorb all the arsenic loaded, showing a lower adsorption capacity. This fact is very likely due to the lower specific surface area of hematite compared to the other adsorbents. Accordingly, hematite showed the highest arsenic uptake per surface area.  63  4.2 ELECTROACOUSTIC TESTS (E.T.) 4.2.1 Magnetite The results for E.T. tests are shown in Figure 4.1, as ESA vs. pH. The ESA data for magnetite in the presence of arsenic (V) and magnetite alone are shown for comparison. The ESA response of magnetite in the presence of arsenic (V) was considerably altered from that of magnetite alone. A comparison of the curves (magnetite and magnetite plus arsenic) shows that this deviation is significant at acidic pH. In addition, both curves show hysteresis that is more pronounced in the presence of As (V). The figure shows that the surface charge of magnetite alone decreases as pH increases. Since the surface charge is positive at pH< 6.5, and negative at pH>6.5, the isoelectric point (I.E.P.) is about pH 6.5. As a consequence of arsenic adsorption, the magnetite surface charge becomes negative over the entire pH range. Also the pH corresponding to the pzc decreases. For magnetite with 50 ppm of As (V), the pzc is located at pH<3.5. However, the reduction of surface charge does not seem to be equally distributed. The shift of the surface charge is less pronounced at higher pH, i.e. a comparison of the curves for both magnetite alone and magnetite with arsenic at pH 10 show similar ESA values. Thus, the decrease of the surface charge as a consequence of arsenic adsorption is less significant at higher pH.  64  Fig 4.1 ESA vs. pHof Magnetite (M) slurries (12.5 g M. /250g slurry, backward titrations are shown as dark points) 4.2.2 Hematite The results of the E.T. with hematite are shown in Figure 4.2. The ESA data for hematite in the presence of different concentrations of arsenic (V) (20 ppm and 50 ppm) and hematite alone are shown for comparison. The ESA profile of hematite in the presence of arsenic (V) was shifted towards more acidic pH values. Comparison of the curves (hematite and hematite plus arsenic at 20 and 50 ppm) shows that this change is less significant at high pH (i.e. pH = 10).  65  -0.8  J  Fig 4.2 ESA vs. pH of HematitefH) Slurries (25gH/250gslurry) The figure shows that the surface charge of hematite alone is decreasing while pH is increasing. The surface charge is positive at pH< 6.5, and negative at pH>6.5. The I.E.P. is placed at a pH of about 6.5. As a consequence of arsenic adsorption, the surface charge of hematite decreases over the entire pH range. Also, the pzc value becomes lower and lower as the concentration of As increases. The pzc decreases from 6.5 to 5.5 at 20 ppm of As (V), and to 5.0 at 50ppm of As(V). As in the case of magnetite, the lowering of surface charge does not seem to be equally distributed over the whole pH range. The shift in the surface charge is less pronounced at higher pH, i.e. the ESA values are similar at pH 10.0. Thus, the decrease  66  of the surface charge towards more negative values as a result of arsenic adsorption is again less significant at high pH. Figure 4.3 presents the ESA data as a function of pH for hematite in the presence of arsenic (V) (20 ppm), and hematite alone, and for the backward and forward titration curves.  -0.8 -I  Fig 4.3 ESA vs. pH of Hematite (H) Slurries (25gH/250g), Back and Forward Titration (backward titration are shown as shadowed points) There is a significant difference between the backward and forward titrations for hematite alone, and the ESA response of hematite shows a clear hysteresis. Moreover, after arsenic adsorption this hysteresis is more pronounced.  67  The results of titration of hematite with arsenic (V) solution, at pH = 4.2 and pH = 9.2, are presented in Figure 4.4  -1 J  1  As(V) solution added, [ml] Fig 4.4 Titration of Hematite Slurry (25 g H/250g) with Arsenic (V) Solution, 5g/L The figure shows a steady decrease of positive surface charge on hematite at pH = 4.2, as well as an increase of the hematite negative surface charge at pH = 9.2 towards more negative values. Both are a consequence of arsenic adsorption. Hence, titration with arsenic solution causes hematite surface charge to decrease even at pH 9.2, when the surface charge is initially negative.  68  4.2.3 Goethite The results for the E.T. are shown in Figure 4.5. The data for goethite in the presence of arsenic (V) and goethite alone are shown together. The ESA profile of goethite in the presence of arsenic (V) was considerably changed from that of goethite alone. In contrast to magnetite and hematite, both curves do not show appreciable hysteresis.  Fig. 4.5 ESA vs. pH of Goethite (G) Slurries (12.5 g G/250g slurry) The figure shows that the surface charge of goethite goes through a point of zero charge at a pH value of about 8.5. There is only small difference between the backward and forward titration, hence hysteresis is not apparent for goethite.  69  As a consequence of arsenic adsorption, the surface charge decreases over the entire pH range. The pzc is situated at pH 6.0 in the presence of 50 ppm of arsenic. In addition, similarly to goethite alone, there is no appreciable hysteresis for goethite in the presence of arsenic (V).  4.2.4 Laterites The electroacoustic data are shown in Figure 4.6 as ESA vs. pH. The data with and without arsenic are shown together for comparison.  Fig 4.6 ESA vs. pH ofLateriteM Slurry (12.5 g L./250g slurry) The figure shows that the surface charge of laterite decreases as pH increases. The IEP is at a pH value of about 7.5. There is a negligible difference between the backward and forward titration, hence, hysteresis is not substantial for the laterite.  70  After arsenic addition, the surface charge curve is shifted towards more acidic pH values with the pzc located at pH 6.5. As in the case of laterite alone there is no appreciable hysteresis in the presence of arsenic (V). The results of the electroacoustic tests for Laterite 2 are shown in Figure 4.7. The ESA data for Laterite 2 in the presence of arsenic (V) and Laterite 2 alone are shown for comparison. laterite - • - Laterite+50ppmAs(V) Laterite+IOOpmAsjV) - 0 - Laterite+200ppmAs(V) Laterite+250ppmAs(V)  Rl  -0.5 J  Fig 4.7 ESA vs. pH of Laterite Slurries (12.5gL/250gslurry) The ESA profile of laterite in the presence of arsenic (V) moves toward lower and lower pH values as the concentration of As (V) increases from 0 to 250ppm. At the same time, the point of zero charge for the laterite sample changes from pH 6.7 to pH<4.0.  71  There seems to be only a minor change in the surface charge and in the pzc values as the As concentration increases from 200 to 250ppm, and in the presence of high doses of As (V) (more than 50ppm). ESA values are quite similar for different arsenic concentrations at high pH.  4.3 ADSORPTION & LEACHING TESTS Adsorption tests were performed prior to leaching tests. Subsequent to adsorption, these samples were leached by MgCh in order to asses the stability of the adsorption products. With the purpose of establish a valid comparison between leaching/ adsorption and electroacoustic tests, the initial arsenic concentration/slurry concentration ratio was the same for all the tests. The following sections will present the results obtained in both tests.  4.3.1 Adsorption Test Adsorption test was performed at pH 5.2, adding arsenic to the corresponding Fe-Ox slurry. It is expected that not all the arsenic added was adsorbed in all the samples. Therefore, after adsorption test some of the samples will likely have some arsenic left in solution. Following the adsorption, each one of the Fe-Ox slurry was split into two equal sub-samples (A & B) in order to asses the arsenic in solution. Samples A and B for the same Fe-Ox had the same quantity of arsenic adsorbed and arsenic in solution. Samples A with no further treatment were submitted to an analytical lab for dissolved arsenic analysis and leaching tests were performed on Samples B (refer to the subsequent section for a description on leaching tests). The analyses of Samples A revealed the  72  amount of arsenic remaining in solution after adsorption (i.e. arsenic in equilibrium or arsenic not adsorbed). Results from samples A for the Fe-Ox studied are illustrated in Table 4.2 Table 4.2 Adsorption Tests Results As adsorbed  As adsorbed  As adsorbed  As in equilibrium.  mg  As adsorbed %  mg/g  mg/m  mg  Hematite(H)  0.752  75.2%  0.752  0.20  0.248  Magnetite(M)  0.964  96.4%  0.964  0.15  0.036  Goethite(G)  1.000  100%  1.000  0.083  0.000  Laterite(L)  1.000  100%  1.000  0.012  0.000  Sample  2  Table 4.2 demonstrates that while goethite and laterite adsorbed all the arsenic, hematite and magnetite had arsenic remaining in solution. Samples A are equal to samples B, thus prior to leaching tests, hematite & magnetite-Samples B have arsenic in solution, while goethite and laterite- Samples B do not have any arsenic left.  4.3.2 Leaching Test Samples B were used in the sequential extraction tests with MgC^. After leaching steps ( l L to 3 L), and washing (W) the filtrated samples were sent to an analytical lab for st  rd  dissolved arsenic analysis. The results are shown in Table 4.3 and 4.4. Since some of the samples B (i.e. hematite and goethite) already contained residual arsenic in solution, the results from the analysis of the first leaching in Column3, Table 4.3 (1 L, C#3) will reflect the arsenic actually leached plus the residual arsenic after st  adsorption.  73  As after first leaching (C#3) - As actually leached + As in equilibrium remaining after adsorption (C# 1)  Eq. 4.1  Thus, the actual amount of arsenic leached in the first step is shown in Column7, Table 4.3(C#7, actual 1 L) and is obtained as the difference between arsenic after first leaching st  in Column3 (C #3) and As in equilibrium remaining after adsorption in Columnl (C# 1). The total arsenic released (R) during leaching is shown by Column8 (C#8), that is the sum of arsenic released from each leaching step: actual first Leaching, Column 7 (Actual 1 L, C#7), second Leaching, Column 4 (2 L, C#4), third Leaching, Column 5 (3 L st  nd  rd  ,C#5) and washing step, Column 6 (W, C#6)). These figures represent the arsenic weakly bound to the Fe-Ox. The fraction of arsenic leached from the mineral surface (i.e. weakly adsorbed arsenic) is given as a percentage in column 9 (R %, C#9). The percent of weakly adsorbed arsenic is equal to the ratio between the total arsenic released from leaching, Column 8 and the arsenic adsorbed, Column 2, multiplied by 100, as demonstrated in the equation below. %Weakly adsorbed As = [Total As released from leaching(C #8) / As adsorbed (A, C#2)] X 100.  Eq. 4.2  The strongly bound arsenic, or arsenic fixed (F %), is shown in column 10 (C #10). And it is calculated by subtracting the difference between 100 and the percent of arsenic weakly adsorbed in Column 9 (C#9).  74  Table 4.3 Leaching Tests Results C#1  C#2  C#3  In Equil. mg As  H  0.248  C#4  C#5  A. mg As  1 L mg As  L mg As  3* L MgAs  Actual1 W L mg As mg As  R mg As  0.752  0.388  0.100  0.058  0.002  0.140  0.299 39.8  60.2  M  0.036 0.964  0.135  0.075  0.054  0.009  0.100  0.237 24.6  75.4  G  0.000  1.000  0.094  0.065  0.059  0.001  0.094  0.218  21.8  78.2  L  0.000  1.000  0.044  0.045  0.047  0.002  0.044  0.138  13.8  86.2  Fe-Ox  st  C#6  C#7  C#8  C#9  C#10  st  R F %As %As  Table 4.4 Leaching Tests Results in a Surface Area Basis Fe-Ox  As Released  As Released  As Fixed  As Fixed  mgAs/g solid  mgAs/m solid  mgAs/g solid  mgAs/m solid  H  0.398  0.10  0.602  0.16  . M  0.246  0.037  0.754  0.11  G  0.218  0.018  0.782  0.065  L  0.138  0.0017  0.862  0.011  2  4.4 ADSORPTION ISOTHERMS Adsorption isotherm tests were performed on samples of Laterite#2 at pH 5.2 and 9.2. Table 4.5, and Fig 4.8 present the results of these tests at pH 5.2.  75  Table 4.5 Arsenic Adsorption onto Laterite at pH 5.2 Sample#  As initial mgAs/L 100.0 100.0 100.0 125.0 100.00  4-2 4-3  4-4 4-5 4-6  As initial mgAs/gsol 1.96 0.98 1.21 4.00 10.00  As in Eq. mgAs/L 0.009 0 0 0.050 13.687  As adsorbed mg/g 1.96 0.98 1.21 4.00 8.63  mgAs/g  0  2  4  6  8  10  12  mg/L As  Fig. 4.8 Adsorption Isotherm atpH 5.2 ( mgAs/g solid vs. mg/L ofAs in equilibrium) Table 4.6, Fig 4.9 and 4.10 present results from adsorption tests at pH 9.2. Fig 4.9 shows adsorption isotherm at pH 9.2, while Fig. 4.10 shows the Langmuir Isotherm Model at this pH.  76  Table 4.6 Arsenic Adsorption onto Laterite at pH 9.2 Sample#  As initial mgAs/L  As initial mgAs/gsol  As final ppm  As adsorbed mg/g  9-2  99.4  2.00  0.1874  2.00  9-3  101.73  0.97  0.0452  0.9.7  9-4  125.0  1.25  0.0701  1.25  9-5  100.0  4.00  4.3820  3.82  9-6  100.0  10.00  37.1180  6.29  7.5  -  mg As/ g O  5  2.5  .  .  0  10  20  30  40  mg/L As  Fig. 4.9 Adsorption Isotherm atpH9.2 ( mgAs/g solid vs. mg/L ofAs in equilibrium)  11  mg/L/mg/g  8  6• 4• 2• n, 0  10  .20  30  40 mg/L As  Fig. 4.10 Langmuir Isotherm Model at pH 9.2 Fig 4.8 & 4.9 show adsorption isotherms of arsenic (V) on laterite. The high affinity of arsenate to the laterite surface is clearly demonstrated by the initial steep slope of the isotherm at low arsenic concentration at both pH 5.2 (Fig. 4.8) and pH 9.2 (Fig 4.9). The maximum adsorption capacity was estimated by fitting a Langmuir-type equation to the adsorption isotherms (Fig 4.10). Applying this model gives a maximum adsorption capacity of 6.41mgAs/g (0.08 mg/m ) on laterite at pH 9.2, and a maximum adsorption 2  capacity of 8.65 mgAs/g (0.11 mg/m ) on laterite at pH 5.2 2  4.5 EVALUATION OF A SIMPLE WATER TREATMENT PROCESS Water samples (lOOmL, 20 ppm of As) were mixed with 5 g of laterite, filtered and then analyzed for arsenic in solution. Table 4.6 presents the results of this method, (mixing time vs. arsenic in mg/L).  78  Table 4.7 Evaluation of a Simple Water Treatment Process Sample #  Time  Final As in solution ppm  T1  15 sec  0.604  12  60sec  0.472  T3  10 min  0.001  T4  1hr  0.001  The tests showed a quick arsenic uptake, indicating a strong affinity of arsenic towards laterite. After 10 minutes of stirring arsenic remaining in solution was at lOppb, and remained unchanged over a further 50min, indicating that adsorption equilibrium is achieved within the first 10 minutes of mixing. The equilibrium arsenic concentration is below the national (Canada) and international limits for arsenic in drinking water.  4.6 DISCUSSION 4.6.1 Electroacoustic Measurements (E.T.) -Surface charge and pH of pzc for Fe-Ox Results from E.M. showed that surface charge is positive when pH<pzc and negative when pH >pzc, and pH of pzc are placed between 6.5 and 8.5 for all Fe-Ox studied. The charge on the oxide surface is established by adsorption or desorption of protons bound to the surface hydroxyl groups. These reactions can be represented by EFeOH <r> E FeOH + H +  2  EFeOH <-> EFeO" + H  +  +  Equation (2.5) Equation (2.6)  79  These equilibrium reactions are pH dependent. An increase of [H+] will be followed by a displacement of the equilibrium to the left, thus at lower, acid pH, neutral and positive groups will predominate and the oxide surface will be positively charged. On the other hand, a decrease of [H+], will be followed by a displacement of the equilibrium to the right, thus at higher, basic pH, neutral and negative groups will predominate and the oxide surface will be negatively charged. In addition, positive, negative and neutral functional groups can coexist on the oxide surface at any pH. At pH < pzc the FeOH2  +  groups predominate over the FeO" groups, i.e. although the surface has a net positive charge some FeO" groups are still present. At the pzc, the number of FeOH2 groups +  equals the number of FeO" groups, and as the pH increases, the number of FeO" groups increases, making the overall surface charge negative. - pzc Shifts as consequence of arsenic adsorption ET tests showed that shifts of pzc as well as a decrease of surface charge occur due to arsenic adsorption. Shifts of the pH of pzc are a very good indication that there is a change in the surface components of the Fe-Ox and that inner sphere complexes are formed. Since ESA is a measure of the particle surface charge, which in turn reflects the proportion of positively, negatively and neutral groups on the surface, a decrease of the surface charge after arsenic addition, indicates that negative arsenate groups are adsorbed onto the particle surface. Adsorption of arsenate (ASO4H2", ASO4H" ) on the Fe-Ox 2  surfaces can take place via three mechanisms: inner sphere surface complex (specific adsorption), outer sphere surface complex, and adsorption in the diffuse-ion swarm (Stumm W., 1992).  80  The formation of an inner sphere surface complex (specific adsorption) during arsenic adsorption onto Fe-Ox is indicated by the decrease of the pH of the point of zero charge. This fact indicates that the anion is bound directly to the particle surface, and because of that the net surface charge decreases or becomes more negative. Thus, the point of zero charge, pzc, is shifted to lower pH values. In addition, the decrease of surface charge after arsenic adsorption implies that adsorption occurs over the whole pH range, even when the Fe-Ox surface charge is negative (pH>pzc). This observation is also evidence that the arsenic anion is bound as inner sphere complex (specific adsorption) onto Fe-Ox surfaces. Specific adsorption may take place on a neutral surface, or even on one with the same charge as the adsorbing species, while non-specific adsorption (outer sphere complex) is dominated by electrostatic forces and thus it occurs only between oppositely-charged species. Thus, for non-specifically adsorbing anions, adsorption is negligible above the pzc when the mineral surface is negatively charged. When electrostatics govern adsorption (i.e. non specific adsorption), the surface must have an overall positive charge in order for anion adsorption to take place, hence the region of maximum adsorption lies in the pH range <pzc of the solid. In contrast, where specific adsorption is involved, an overall positive surface charge is not required (only FeOH2 and FeOH groups) which explains why adsorption can occur at +  pH above the pzc (Cornell and Schwertmann, 2003). -Reversibility and hyteresis of ET curves Goethite and laterite showed reversibility of ET curves, indicating adsorption as the predominant mechanism for arsenic bound to the surface of Fe-Ox. Furthermore, magnetite and hematite E.T. curves showed hysteresis. This fact may suggest that the  81  surface reaction between arsenate and magnetite/hematite changes from adsorption into surface precipitation. In addition, hematite showed a positive hysteresis, its surface charge increased with time, while magnetite showed a negative hysteresis, its surface charge decreased with time. Dzombak & Morel (1990) stated that hysteresis of electrokinetic curves may imply some surface precipitation or it may be the consequence of the kinetics of the reaction of adsorption / desorption. They describe surface precipitation as the formation of a different solid phase, a solid solution whose formation starts when the - saturation concentration is exceeded. Alternatively, anion precipitation can be viewed as a ternary adsorption (multilayer adsorption), starting when metal ions (Fe) adsorb onto the adsorbed anion (arsenate). The two models are not in conflict, though multilayer adsorption does not require that the solution be saturated. Furthermore, the terms ternary adsorption, polymer formation on the surface and surface crystal formation are used as equivalent of surface precipitation (Ler & Stanforth, 2003). In order for surface precipitation to occur the Fe-Ox must dissolve to contribute iron to the surface precipitate. Fe-Ox have low solubility in water, but the solubility increases at acid pH, especially at the lowest pH used in the study (pH 3). The iron dissolved at such low pH may be sufficient to form the surface precipitate, as reported by Ler & Stanforth (2003) in the case of phosphate precipitated on goethite. The iron product of this dissolution is adsorbed onto the surface, and then new arsenate may be added. Thus as a consequence the surface charge will change over time. Changes (increase) in surface charge over time as indication that surface precipitation is occurring are reported by Ler & Stanforth (2003).  82  The change of surface charge over time showed by the ET (hysteresis), may suggest that surface precipitation is taking place onto the hematite and magnetite surfaces. Therefore, the surface reaction between arsenate and magnetite and arsenate and hematite may change from adsorption into surface precipitation, meaning the formation of a ternary complex (multilayer coverage).  4.6.2 Adsorption Tests Results from the tests performed prior to leaching showed that arsenic was totally adsorbed onto goethite and laterite, while some arsenic remained in solution from tests with hematite and magnetite. Goethite and laterite did not reach their saturation capacities therefore they adsorbed. 100% of arsenic loaded. Meanwhile, hematite and magnetite reached their saturation point, hematite adsorbed 75.2% of arsenic, and magnetite 96.4% of arsenic (Table 4.2). -Saturation and hysteresis From the adsorption tests, hematite left 25% of arsenic in solution, while magnetite only 4%. The results imply that both surfaces were saturated. In addition both minerals presented hysteresis in their ET curves, suggesting either precipitation or kinetic causes. According to Sadiq (1995), precipitation of an arsenic solid phase on soils may occur on soil colloid surfaces subsequent to its adsorption. In that case the adsorbed anion acts as a sorption site for dissolved iron, forming a ternary complex (or surface precipitate). Meanwhile the bulk solution does not need to be saturated or reach a given concentration of either anion or cation in solution, as it has been shown in previous studies of cation adsorption/precipitation (Ler and Stanford, 2003). Thus, precipitation takes place after  83  adsorption sites are occupied, when the surface mineral is saturated with respect to the anion. Hematite and magnetite both show saturation of the surface mineral suggesting that they may exhibit surface precipitation.  4.6.3 Leaching Tests Leaching tests were performed to target the ionically (weakly) bound arsenic, assuming that arsenic that is not ionically bound is either strongly adsorbed or coprecipitated with Fe-Ox (Keon et.al, 2001). The results from leaching tests showed a correlation between the quantity of arsenic strongly bonded (% weight) and the surface area of the Fe-Ox mineral i.e arsenic strongly bonded increases while surface area increases. In addition, results support assumptions that specific adsorption and precipitation are the two prevailing mechanisms for arsenic adsorption onto hematite and magnetite, while specific adsorption is the one for goethite and laterite. -Arsenic strongly bonded and surface area Results from leaching tests suggest a correlation between the specific surface area of the Fe-Ox minerals and the quantity of arsenic fixed (% in weight, Table 4.3, Column 10). Laterite showed the highest proportion of arsenic strongly bound (86.2%), then goethite (78.2%), magnetite (75.4%), and finally hematite (60.2%). The correlation is not linear and showed that arsenic fixed approaches a constant value while the surface area of FeOx increases.  84  Fig. 411 As Fixed% vs. Fe-Ox Surface Area -Binding ofArsenic to the Fe-Ox surfaces Arsenic not leached by this test is either strongly adsorbed or coprecipitated with Fe-Ox (Keon et.al, 2001). The leaching tests results showed high proportion of arsenic not ionically bound which is strong proof of arsenic inner sphere complex formation/ precipitation onto Fe-Ox. (Tables 4.3 & 4.4) The highest quantity of arsenic strongly bonded or fixed on a surface area basis (Table 4.4) is shown by hematite (0.16mgAs/m ), followed by magnetite (O.llmgAs/ m ), then 2  2  goethite (0.065mgAs/ m ) and finally by laterite (O.OllmgAs/ m ). The fact that hematite 2  2  and magnetite showed the highest value of arsenic fixed per area is in agreement with the assumption that precipitation plus specific adsorption occur over their surface, while goethite and laterite bound arsenic only through specific adsorption, thus they yield a lower quantity of arsenic strongly bound. The results correlate well with the assumption of ternary complex formation (precipitation) plus inner sphere complex formation (specific adsorption) for hematite and magnetite and inner sphere complex formation for laterite and goethite. Precipitation and specific adsorption imply tightly bound products. Thus, arsenic adsorbed onto these Fe-Ox forms stable surface products.  85  4.6.4 Adsorption Isotherms Adsorption isotherms (graphs 4.8 to 4.10) showed that arsenic adsorbs on laterite at both pH 9.2 and 5.2. However, laterite showed the highest As(V) adsorption at pH 5.2. There is a decrease of arsenic adsorption with pH, which is shown by the lower adsorption at pH 9.2. The results indicate that pH 5.2 is favorable for arsenic adsorption. This fact is also shown by the maximum adsorption capacity, i.e. 8.65 mgAs/g, 0.11mg/m of 2  laterite at pH 5.2, and 6.41mgAs/g, 0.08 mg/m of laterite at pH 9.2. Furthermore, the 2  shape of the isotherms suggests a high relative affinity of arsenic towards the laterite surface (Sposito, 1984). -Adsorption and ET curves Fig 4.12 presents the results from the electroacoustics measurements and the adsorption tests at pH 9.2 and 5.2 for arsenic(V) onto laterite. Data for this graph come from Tables 4.5 & 4.6 and Fig 4.7. The graph shows arsenic adsorbed (as mmolAs/g solid) at pH 9.2 and 5.2, and the corresponding ESA measurement; both are plotted vs. As initial concentration in mg/g. There are three curves for arsenic adsorption, one straight line is the complete adsorption situation where all the arsenic added is totally adsorbed, and two curves corresponding to the actual arsenic adsorbed at pH 9.2 and pH 5.2. ESA curves at pH 9.2 and pH 5.2 show variation of surface charge with As.  86  0  2  4  As Initial, mgAs/g  6  8  10  12  Fig.4.12 a) As Adsorbed (mmol As/g laterite x 10) vs. As Initial (mgAs/g laterite) and b)Surface Charge (ESA) vs. As Initial (mgAs/g laterite) Adsorption of arsenic anion on laterite correlates well with the ET curves-both of them indicate that adsorption occurs at both pH acid (5.2) and basic (9.2). The graph shows that an increase in arsenic adsorption at pH 5.2 is followed by a corresponding decrease of surface area charge. Also the final declining in arsenic adsorption at pH 9.2 is accompanied by the tendency for the surface charge to approach a constant value.  4.6.5 Evaluation of a simple water treatment process A simple method was developed to remove arsenic from water using natural iron mineral bearing soils. Lateritic soil was used because it is naturally abundant, an efficient adsorbent for arsenic and has a good final product stability. The method involves mixing laterite with arsenic contaminated water (20mg/L) for ten minutes and then filtering, using a coffee filter.  87  The treatment lowered arsenic levels in the treated water below drinking water standards. (lppb< 25ppb, Canadian Drinking Water Standard) The treatment method is inexpensive, and simple, making it suitable for household use. For example, a family of five members has water needs for drinking and food preparation of about 12.5 L per day. Thus a sack of 100kg of this laterite will treat enough water to meet the family needs for 160 days (2,000L).  4.6.6 Comparison between Fe-Ox Hematite, magnetite, goethite and laterite have been studied in their role as arsenic adsorbents. Results showed that all of them are suitable as arsenic adsorbents. ET tests showed that arsenic adsorption occurs over the whole pH range considered (4-11) and also that Fe-Ox have I.E.P. between 6.5 and 8.5. The adsorption capacity of laterite was estimated to be 0.11 mg/m , while magnetite 2  9  9  adsorption capacity was 0.15 mg/m , and hematite 0.20mg/m . Leaching tests by MgCL; were performed to study the stability of the adsorption products and results expressed on a weight percentage basis showed that hematite had 60.2%, magnetite 75.4%, goethite 78.0% and laterite 86.2% of arsenic strongly fixed. While these results expressed on a surface area basis showed that hematite had 0.16mg/m 9  9  hematite, magnetite O.llmg/m, goethite 0.065, and laterite 0.011 mg/m of arsenic strongly fixed (Tables 4.3 and 4.4). These results showed that most of the arsenic is strongly bonded to the Fe-Ox surface (i.e. specific adsorption or precipitation) they indicate that the surface products formed are stable.  88  5. CONCLUSIONS & RECOMMENDATIONS  5.1 CONCLUSIONS Research general objective has been fulfilled; iron oxides convenience and potential for arsenic removal in water have been shown. Research questions, previously developed in chapter one, have these answers: 1) Electrokinetic and adsorption tests have shown the capability of various Fe-Ox for arsenic adsorption in a broad pH range (4-11). 2) Arsenic adsorption takes over the whole pH range (4-11), though higher adsorption capacity have been shown at pH <pzc (point of zero charge). 4) Shifts in pzc of Fe-Ox can be used as evidence of strong specific ion adsorption and inner-sphere surface complex formation; such shifts have been observed following arsenate adsorption on the various Fe-Ox studied. 5) Arsenic adsorption onto Fe-Ox products shows a high grade of stability. Results from leaching tests corroborate the inner sphere complex formation assumption from electrokinetic tests. Most of arsenic is attached to the solid through specific adsorption. This fact collaborate with the understanding about potential mobility of arsenic from the solid phase, for example, arsenic that is not ionically bounded can not be mobilized by an increase in the ionic strength of the aqueous phase, such as increase of salinity of the water, thus is more stable product. 6) Laterite looks as a promising tool for arsenic removal from water. Advantages of using a natural lateritic soil for arsenic removal have been shown. Naturally occurring FeOx are attractive for arsenic removal from contaminated water, because they are cost  89  effective and their physico - chemical properties make them suitable for arsenic removal over a wide pH range. Moreover, laterite as common, abundant natural iron mineral bearing soil is a promising tool for arsenic removal because it is cheap, abundant, efficient adsorbent for arsenic, and has good final product stability. 7) A simple method for arsenic removal using laterite was developed. The method involves mixing laterite with arsenic contaminated water for ten minutes and then filtering, using a coffee filter. After treatment, arsenic levels in the treated water were below drinking water standards. (lppb< 25ppb, Canadian Drinking Water Standard) The method is inexpensive, and simple, making it suitable for household use. Moreover, There is not need of pH adjustment for this treatment, the adsorption reaction is fast and the treatment device is simple. 8) Electroacoustic tests measurement techniques are simple, and because they do not involve phase separation, less problematic. They can be done at high solid content. And, because of the strong influence of pH, sorption data need to be presented as function of pH, so electrokinetic results bring a broad view of adsorption vs. pH. Furthermore, information about adsorption mechanisms is also given by electrokinetic tests because of shift of pzc.  5.2 RECOMMENDATIONS Recommendations derived from the main topics of this research are presented below. Based on the results of experiments, subsequent research on this topic should address the following issues 1) Extend this study to include more natural minerals and soil bearing minerals.  90  2) Extend development of the household method to apply laboratory results to case study parameters. 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"Environmental Health Criteria 224: Arsenic and Arsenic Compounds". http://www.inchem.Org/documents/ehc/ehc/ehc224.htm#4.0 Zhang W, Singh P, Paling E, Delides S (2004) " Arsenic Removal from Contaminated Water by Natural Iron Ores" Minerals Engineering (2004) 17: 517 -524  100  APPENDIX  101  Technical Data PtaUBH 1,5  Stack Iron Oxide  "'  • *-  Etemenlts Synthetic Black Iron Qjdctea ara precipitated magnetite (Fa approximately 24% expressed as F « 0 .  jQ«X  Typically, the ferrous content to ,  Thay yield daop black masttofte color and rich bluish Mack tints,  hawexcellent hiding pow#r, arte) art low In heavy nutate content Wsateg,  dispersion and susp««*lctfi propsrtlaa are Bueeltarit.  Parttda afza la uniform wilt) an avcr*ge  xtamelBr of appraxirnat afy (Mfntlerwi*ThBfi»caJlert color, lint and higrvhidtng of these product* *ugge$t u&e In palms for metal furniture, Are > aacqpe^graytirrtliMl^a^  Other u w c include prtMbis M o h soncrete products and floprln*:'  ^wtiare pormanBrrt, ew^-wettlng black pigments are desEraWo. For  erflteal applications,  Etemenli*' Pigments high (togman OK 53990 will meet ttie most ojaactlng  riqulremante, EK-5495D I* designed to yield a SHsgrnan grind on high spaed disperaara. TYPICAL PROPERTIES (Bate.BdfeWAppJy To AB Black Grades} Forintilir • . Parflcte.Shap* . : .  FSjO« cubfc  H*gman(Dgredaonly)  6  Mesh RBtenttort, % . :  DH  % PUXfTY  Specific RoBlatanco.Q Sp&eSte Gravity Refractlvs Irxiax Nwmbor  2,42  AVERAGE PARTICLE StZB  SPSORC SURFACE AREA (m>®  <w. 93  0.02 6 SOOC s.rj • .  C.39  :  0.0  Aairifan^i*(*ven:am^^ N o guarantee, eilfter mpeeea&lorlirpl^fem^B^r*^^1h«ateor!««thraapactlsthelrrWngmiwrtoifanypilwTt Oarp»wiuefeanswltf»<lh 1l»underalmdW usertosslsfyr«£pensiMefertfetemtinii^g thatr sultatalrty forara'purpese, J H i N & i i n j & n is m ifaba c h i s e l . .BIMSIS in a v t f e f l c a ^ y e r e a t t g V far p u M c s S a n < r p u M i c ' d l M f e a t o n iiwilhaut written pamnagfaa Hera'anlta Flomenta tod' Vii  9M-tm  EismerrtisP^raentelnc. • 11 tmctiUva Bf«8» Suite I • F i l / v l o w HeiffiW, I i USA talo^or.e: ar^^3.779e * F ^ i r t i i l K Sia.S2B.aQ2S W W site;.wttw.inementisriBTOfih.aotn « e.mailr pamwtsglHl.i^is ,-r..n  102  - MATERIAL 'SAFETYDATA SHEET fiemeuos Pigmwits Fix. 2051 Lynch Avenue BwtSt LoaK.lL S32G4  Pttidua: MSDS Vo.  PKMEKT..' EFM3 i 8 .' Revisions April 20OJ  HAZARD MATERIAL IDENTIFICATION SYSTEM tteatt Based Fiaamiibtlity Hsaud  SECTION!.  0-Miramsl 0 -Minumi 0 -Minimal E - Glasses, Qtovra, Dust Resp  MATERIAL IDENTIFICATION YELLOW IRON OXIDE Yellow Iron Oxede, C J , Pigment Yeitow 42. C.J. Ne, 7 7 4 «  YP-2100.  QantMatNaips: Blsmeirtw Pigments Inc. 2051 Lynch Avenue E a s t S L L a u k l L <j220i  SECTION/!,  PHONE:  61&-64&22QI E . S t Louis Plsrtt  INGREDIENTS AMD HAZ4&DS'  INGREbfEOTMAMBr tfOtl Oxfife Yellow  5I274-0&-I  >99  ACGJHTLV:  SOmg.Trf'TWA  OSHA PEL:  t5 ms/M* TWA  103  Table 4.1.1 Scoping Tests: As adsorbed from different Fe - Ox Slurries Slurry  Solution Final Volume As in equil. As in equil. As adsorbed # Solution corrected (mL) mg/L mg/L mg/L  As-Standard Solution  E1  Hematite  E2  Magnetite  E4  105.4  Goethite  E5  Laterite#1  121.0  3.9585  4.7898  N  3.9211  1.0289  <0.0005  <0.0005  4.9500  120.8  <0.0005  <0.0006  4.9500  E6  122.3  0.0007  0.0009  4.9491  Laterite#2  E7  137.6  <0.0005  <0.0007  4.9500  Blank  E8  100.4  <0.0005  <0.0005  4.9500  standard error 3.3%  102  3.8442  Table 4.1.2 Scoping Tests: As adsorbed from different Fe - Ox Slurries Slurry  Solid As adsorbed As adsorbed As adsorbed Surface Area m /g mg/L mg/g solid mg/m solid 2  2  Hematite  3.77  1.03  2.1E-01  5.5E-02  Magnetite  6.58  4.95  5.0E-02  7.5E-03  Goethite  12.1  4.95  5.0E-02  4.1E-03  Laterite(#1&#2)  81.2  4.95  5.0E-02  6.10E-04  105  ADSORPTION AT Ph 5.2 Sample* # 4-2 4-3 4-4 4-5 4-6  w g of solid  mLAs-added Final vol mL  5.10 10.20 10.30 2.50 1.00  Sample#  w  2.00 2.00 2.50 2.00 2.00  As mg/L mgAs/gsol  100.00 100.00 100.00 100.00 100.00  As- mg/L  100.00 100.00 125.00 100.00 100.00  Asfinal  Q  g of solid  initial  initial  ppm  mg/g  4-2  5.10  100.00  1.96  0.009  1.96  4-3  10.20  100.00  0.98  0  0.98  4-4  10.30  125.00  1.21  0  1.21  4-5  2.50  100.00  4.00  0.050  4.00  4-6  1.00  100.00  10.00  13.687  8.63  As-initial mg 10.00 10.00 12.50 10.00 10.00  mgAs/gsol Asfinalppb Asfinal initial mg/L 1.96 8.9 0.98 < 5 1.21 < 5 4.00 50 10.00 13686.8  0.0089 0.00 0.00 0.05 13.69  mgAs/g Q 1.96 0.98 1.21 4.00 8.63  106  Arsenic adsorption onto laterite at pH 9.2 M Sample* w # g of solid 9-2 9-3 9-4 9.5 ,9.6  v mLAs-added Final vol mL 5.00 10.30 11.00 2.50 1.00  Sample mgAs/gsol initial  2.00 2.00 2.75 2.00 2.00  As final ppm  c0 Asfinal As-mg/L ppb  100.6 98.3 110 100 100  As adsorbed mg/g  9-2  2.00  0.1874  2.00  9-3  0.97  0.0452  0.97  9-4  1.25  0.0701  1.25  9.5  4.00  4.3820  3.82  9.6  10.00  37.1180  6.29  Asfinal mg/L  99.40 187.4 0.1874 101.73 45.2 0.0452 125.00 70.1 0.0701 100.00 4382.4 4.38 100.00 37118 37.12  Q mg/g  1.996229512 0.970442412 1.249299 3.824704 6.2882  mgAs/gsol initial  2 0.9709 1.25 4 10  Langmuir Isotherm at pH 9.2 mg/L mg/g C c/q C q 0.0452 0.97 0.046598 0.0701 1.25 0.05608 0.1874 2 0.0937 4.3282 3.82 1.133037 37.118 6.29 5.901113  0.0452 0.0701 0.1874 4.3282 37.118  C C/q  0.0452 0.046598  0.1874 0.0937  slope correlation  0.15611 0.997531  0.0701 0.05608  4.3282 1.133037  37.118 5.901113  max adsorption capacity  6.405744 mgAs/g sol  108  LEACHING EXPERIMENTS Sample  As(ppm)  GB  blank  Go  not leaching  G1  Sample vol.(L) As(mg)  0.174  0.05  0.0087  2.9  0.0029  0.05  0.000145  first-leach  940.9  0.9409  0.1  0.09409  G2  secondjeach  646.1  0.646  0.1  0.0646  G3  thirdh leach  592  0.592  0.1  0.0592  G4  washing  56.3  0.0056  0.1  0.00056 0.21845 0.000145 0.218305  As released: As in Go Total As released:  1.00 mg  As adsorbed %Asreleased from  21.8%  Goethite  78.2%  As fixed in Goethite  Sample  Asppb  HB  blank  HO  not leaching  H1  As(ppm)  Sample vol.(L) As(mg)  0.0908  0.05  0.00454  4961  4.961  0.05  0.24805  first-leach  3882.3  3.8823  0.1  0.38823  H2  secondjeach  1002.2  1.0022  0.1  0.10022  H3  thirdh leach  577.8  0.5778  0.1  0.05778  H4 As from leaching:  washing  15.8  0.0158  0.1  0.00158 0.54781  Total As released  0.29976  As adsorbed %released As fixed in Hematite  0.75 Hematite  39.9 60.1  Sample  As(ppm)  0.065  Sample vol.(L) As(mg)  LB  blank  L0  not leaching  L1  first-leach  437  0.437  0.1  0.0437  L2  secondjeach  449  0.449  0.1  0.0449  L3  thirdh leach  474.4  0.4744  0.1  0.04744  21.1  0.0211  0.1  0.00211 0.13815  <0.5  L4 washing Total As from leaching: As not leached  0.05 0.05  (  As adsorbed %released As fixed in laterite  0.00325  1.00 Laterite  13.8 86.2  Sample  As(ppm)  Sample vol.(L) As(mg)  MB  blank  225.5  0.2255  0.05  0.011275  M0  not leaching  721.1  0.7211  0.05  0.036055  M1  first-leach  1347.4  1.3474  0.1  0.13474  M2  secondjeach  753.1  0.7531  0.1  0.07531  M3  thirdh leach  536.8  0.5368  0.1  0.05368  92.2  0.0922  0.1  0.00922 0.27295  M4 washing Total As from leaching: As not leached  0.036055  total Asreleased  0.236895  As adsorbed %released  0.96 Magnetite  As fixed in magnetite  24.6 75.4 110  


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