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Kinetic investigation of low-pH Fe(II) oxidation and development of a method for Fe(III) regeneration as part of a process aimed at H2S(g) removal Research report submitted to Grand Water Research Institute Ori Lahav and No'omi Levi May 2009 Project No. 2008724 For the elimination of any doubt, it is hereby stressed that the staff member and/or the Technion and/or the Technion Research and Development Foundation Ltd. Will not be liable for any property damage and/or corporal damage and/or expense and/or loss of any kind or sort that will be caused or may be caused to you or anyone acting on your behalf, in consequence of this statement of opinion or this report, or in any connection to it. Copyright ©: Year of Publication: 2009; By: Ori Lahav and Noomi Levi and the Technion Research and Development Foundation Ltd. 1 Table of contents pp Abstract 4 Abbreviations 7 Symbols 7 1. Introduction 8 1.1. Hydrogen Sulfide (H2S) 8 1.1.1. Properties of H2S(g) 8 1.1.2. Origins and concentrations of H2S(g) 8 1.1.3. Adverse effects of H2S(g) 10 1.1.4. H2S(g) threshold concentration 12 1.2. Methods for H2S(g) removal 13 1.2.1. Gaseous phase and solid phase processes 14 1.2.2. Physical-chemical processes in the aqueous phase 15 1.2.2.1. Regenerative gas scrubbing 18 1.2.2.2. LRSR process 18 1.2.3. Biological processes in the Aqueous phase 20 1.2.3.1. Biological methods 20 1.2.3.2. Microorganisms of particular interest for H2S(g) removal 21 1.2.4. The Fe(III)/Fe(II) LRSR process 22 1.2.4.1. Execution of LRSR processes at neutral to moderately high pH levels 23 1.2.4.2. Execution of LRSR processes at low pH levels 1.3. The kinetics of Ferrous iron (Fe(II)) oxidation by oxygen 26 34 1.3.1. pH dependence 34 1.3.2. Fe(II) oxidation by O2 in natural waters 39 1.3.3. Fe(II) oxidation by O2 at low pH levels 41 1.4. Electrochemical oxidation of Fe(II) 44 1.4.1. Direct oxidation of Fe(II) 44 1.4.2. Indirect oxidation of Fe(II) 44 1.4.2.1. Competing reactions 45 1.4.2.2. Oxidation of Fe(II) by Cl2 46 2 Contents (continued) 1.4.3. Electrolytic production processes of active chlorine 46 1.4.3.1. Divided or undivided cell 46 1.4.3.2. Batch or flow-through mode 47 2. Hypothesis and objectives 50 2.1. Research hypothesis 50 2.2. Research objectives 51 3. Materials and Methods 52 3.1. Chemicals 52 3.2. Analytical equipment 53 3.3. Experimental 53 3.3.1. H2S(g) reactive-absorption experiments 53 3.3.2. Determination of practical precipitation potential of the working solutions 54 3.3.3. Determination of Fe(II) oxidation rate 54 3.3.3.1. Catalytic oxidation of Fe(II) 54 3.3.3.2. Electrochemical oxidation of Fe(II) 55 3.4. Analytical methods 57 3.4.1. Determination of the total dissolved iron, phosphate and copper concentrations 57 3.4.2. Determination of dissolved ferrous iron concentration 57 3.4.3. Determination of dissolved ferric iron concentration 57 3.4.4. Determination of chloride concentration 57 3.4.5. Determination of ammonia concentration 57 3.4.6. Analysis of species distribution by the MINEQL+ software 57 4. Results and Discussion 60 4.1. Catalytic oxidation 60 4.1.1. H2S(g) reactive-absorption efficiency 60 4.1.2. Practical precipitation potential in the working solution 62 4.1.3. Fe(II) oxidation rate in the catalytic oxidation experiments 64 4.1.3.1. Effect of pH 65 4.1.3.2. Effect of initial concentration of Fe(III) 68 3 Contents (continued) 4.1.3.3. Effect of initial concentration of Cu(II) 69 4.1.3.4. Effect of total phosphate concentration 73 4.1.3.5. Effect of total sulfate concentration 75 4.2. Electrochemical Fe(II) oxidation as part of the LRSR process 79 4.2.1. H2S(g) reactive-absorption efficiency 79 4.2.2. Practical precipitation potential in the working solution 83 4.2.3. Fe(II) oxidation rate in the electrochemical oxidation process 84 4.2.3.1. Direct electrooxidation 85 4.2.3.2. Indirect electrooxidation - effect of anode to cathode surface area ratio (Sa:Sc) 85 4.2.3.3. Indirect electrooxidation - effect of current density 87 4.2.3.4. Indirect electrooxidation - effect of chloride concentration 89 4.2.3.5. Potential chlorine loss 91 4.2.3.6. Indirect electrooxidation - energy cost 91 4.2.3.7. Volume of solution in the electrolytic reactor 91 4.2.3.8. Electrode material 92 4.2.3.9. Changes in pH 92 5. Conclusions 93 Reference list 94 4 Abstract Hydrogen sulfide (H2S) is a product of microbiological anaerobic activity. H2S is emitted from environmental facilities and industrial processes such as petroleum refineries, paper manufacturing, anaerobic digestion processes etc. The removal of H2S from gaseous streams is required for the health of the general public, occupational safety and for operational reasons. The Liquid Redox Sulfur Recovery (LRSR) process is a common and promising process for H2S(g) removal that is based on reactive absorption (i.e. absorption of the gas into an aqueous solution followed by a chemical reaction). The H2S(g) is absorbed into solution and then oxidized to S0 by an intermediate redox couple. The most currently used couple is Fe(III)/Fe(II). Ferric (Fe(III)) oxidizes H2S(aq) to S0, and the ferrous (Fe(II)) formed is oxidized back by oxygen to Fe(III). Working at pH7 to pH9 is favorable since the absorption of H2S(g) is efficient and the spontaneous oxidation of Fe(II) by O2 is very rapid. However, the addition of organic chelates is essential in order to avoid rapid precipitation of Fe(III) species as ferric oxides or oxy-hydroxides. At low pH levels (pH~2) the solubility of ferric species is high enough to avoid the need for chelating agents. Since the rate of Fe(II) oxidation is very low at this pH range, bacteria populations have been suggested to catalyze the regeneration process, mainly from the group Acidithiobacillus Ferrooxidans. However, this method suffers from two major drawbacks: (1) the dependence of the process on the sensitive autotrophic biomass and (2) the relatively rapid precipitation of Fe(III) solids, mostly of the jarosite group. The current work investigates two new approaches for applying the LRSR process at pH1.0. The first method is based on catalytic oxidation of Fe(II) by O2 in the presence of copper and phosphate, which are known to have a catalytic effect. The second is based on electrochemical oxidation of Fe(II), directly (oxidation of Fe(II) on the anode) or indirectly (oxidation of Fe(II) by chlorine formed on the anode). The very low operational pH (pH1.0) was chosen in order to minimize precipitation of ferric species for long-term LRSR operations. Two other factors to be considered in the process are the efficient reactive-absorption of H2S(g) and rapid oxidation of Fe(II). The catalytic oxidation method, although proven to accelerate Fe(II) kinetics in lab tests, was found to be infeasible for the LRSR process, since despite of the low pH a 5 H2S introduced to the reactor precipitated with copper to form CuS, and Fe(III)phosphate solids also precipitated quite rapidly. The direct electro-oxidation method was also found inapplicable since the efficiency of the reactive-absorption of H2S(g) was very low in the absence of chlorides. In contrast, the indirect electro-oxidation method was found to be highly feasible: the reactive-absorption efficiency of H2S(g) was high in the presence of chlorides, there was no appreciable amount of precipitates and the oxidation rate of Fe(II) was very high. 6 Abbreviations AC Active chlorine AD Anaerobic digestion A.F. Acidithiobacillus Ferrooxidans AFO Animal feeding operations DC Direct current DO Dissolved oxygen FAS Ammonium iron(II) sulfate 6-hydrate GRT Gas retention time LRSR Liquid Redox Sulfur Recovery ORP Oxidation-reduction potential Symbols E Efficiency of reactive-absorption Fe(II)T Total ferrous iron concentration Fe(III)T Total ferric iron concentration Ox The oxidized form of the redox PT Total phosphate concentration R- The reduced form of the redox rFe ( II ) Rate of Fe(II) oxidation rH 2 S( g ) Rate of H2S(aq) oxidation ST Total sulfate concentration 7 6. Introduction The economic prosperity of any country is linked to the level of energy consumption. Throughout the world, the increasing demand for energy is attributed to the growing human population and subsequent desire to improve the living standard through industrialization. The energy requirement for keeping pace with the development is met through either conventional (coal, oil) or non-conventional (air, solar) energy sources. Among those sources are gaseous fuels - combustible gases such as natural gas (fossil origin) and biogas (anthropogenic origin). Gaseous fuels contain hydrogen sulfide (H2S) in significant concentrations, along with other sulfur species, which exist at traces concentrations. Different industries generate gaseous streams that contain appreciable concentration of H2S (Pandey et al., 2004). 6.1. Hydrogen Sulfide (H2S) 6.1.1. Properties of H2S(g) Hydrogen sulfide (H2S) is a flammable, colorless gas that smells like "rotten eggs" at low concentrations. Sulfur containing compounds in general (such as mercaptants) and H2S in particular are the products of the biodegradation of sulfur containing amino acids such as cysteine and methionine by anaerobic bacteria. It is also a product of anaerobic reduction of sulfate (SO42-) (Firer et al., 2008). H2S is soluble in various liquids including water and alcohol (WHO, 2000). Henry’s constant for H2S is 0.102 M·atm-1 (Stumm and Morgan, 1996). In the aqueous phase H2S(g) is a diprotic weak acid that dissociates according to the following equilibrium reactions (Noyola et al., 2006): H 2 S ( aq ) ↔ HS − + H + HS − ↔ S 2 − + H + pKa1 = 7.0 pKa 2 = 12.9 (1) (2) 6.1.2. Origins and concentrations of H2S(g) Only 10% of the total H2S global emission is of non-anthropogenic origin (Malhotra et al., 2002), e.g. volcanic gases and hot springs (Busca and Pistarino, 2003). The remaining is emitted from environmental facilities and industrial processes such as petroleum refineries, coke production, viscose rayon production, the tanning industry, wastewater treatment plants, paper and pulp manufacturing, food processing, 8 extraction of natural gas and anaerobic digestion (AD) processes (WHO, 2000; Malhotra et al., 2002; Busca and Pistarino, 2003; Ramírez-Sáenza et al., 2009). Sodium sulfide (Na2S) is one of the most widely used reagents in the tanning Industry. Processing liquors possess sulfide ion concentrations up to 2000 mg/kg and untreated tannery wastewater contain up to 20 mg/kg sulfide (Lawrence et al., 2000). The content of sulfur in crude oils is typically in the range of 0.3–0.8 wt.%. A systematic increase in the sulfur content of extracted crude oils has been recorded over the past two decades and a further increase is anticipated (Lawrence et al., 2000). For example, Pandey et al. (2004) recorded an H2S(g) concentration to be 52,600 ppm in a refinery fuel gas sample. The composition of typical refinery fuel gas is presented in Table 1. Natural gas is a complex mixture containing desirable gaseous hydrocarbons and nonhydrocarbon components such as H2S, CO2 and water. According to Koros and Mahajan (2000), the world market for natural gas is estimated at approximately 22 billion US$ annually. Around 40% of the approximately 54 trillion standard cubic feet of gas reserves in the lower 48 United States are not being developed due to the high cost of gas treatment (Koros and Mahajan, 2000). The H2S content of natural gas ranges from 100 up to 300,000 ppm and is expected to increase (Lawrence et al., 2000). Anaerobic digestion processes are an effective technology for the reduction of the organic matter and simultaneous production of energy through biogas production. Biogas is a sub-product of anaerobic digestion that has a high energy value attributed to its high methane content. The use of this gas as a fuel can decrease both energy costs and operational costs in waste treatment plants where it is generated (Mesa et al., 2004). It is also used in industrial-scale animal feeding operations (AFOs) to treat animal wastes and attaining the same benefits. Biogas production and utilization is constantly increasing, as it represents a “green”, renewable energy, obtainable in a relatively economical way (Cosoli et al., 2008). Biogas is generally composed of 60–65% methane (CH4) and 35–40% carbon dioxide (CO2). Minor constituents include hydrogen sulfide (H2S), nitrogen gas (N2), hydrogen gas (H2) and traces of oxygen (O2), carbon monoxide (CO), ammonia 9 (NH3), argon (Ar2) and other volatile organic compounds (VOC). The composition of biogas depends on the type and concentration of organic matter to be digested, on the physicochemical conditions in the digester (pH, alkalinity, temperature) and on the presence of other electron acceptor species such as sulfates and nitrates (Noyola et al., 2006). A typical composition of a biogas is given in Table 2. Table 1. Characteristics of a refinery fuel gas sample (Source: Pandey et al., 2004) Component Composition (vol.%) Table 2. Typical composition of biogas (Source: Pauss et al., 1987, quoted by ter Maat et al., 2005) Component Composition (vol.%) CH4 47.76 CH4 52–95 H2S 5.26 H2S 0.001–2 H2 1.61 H2 0.01–2 C2H6 17.87 CO2 9–45 Propane 11.12 N2 0.1–4 Propylene 4.88 O2 0.02–6.5 i-Butate 1.61 Ar 0.001 n-Butane 3.73 CO 0.001–2 Butylene 3.97 NH3 Trace 1.71 Organics Trace Pentane 4 1 vol.% = 10 ppm As can be seen from Table 2 the range of H2S content in biogas vary considerably. Concentrations in a similar range were reported also by Mesa et al. (2004), Chung et al. (2006), Noyola et al. (2006), Qaisar et al. (2007), Cosoli et al. (2008) and Fortunya et al. (2008). Much lower concentrations were measured in AFOs: Heber et al. (2004, quoted by Gendel, 2007) measured H2S concentrations throughout the year in swine facilities and obtained concentrations between 0.14 and 1.20 ppm. Okoli et al. (2004) measured concentration of H2S in five intensive layer farms during the month of August and obtained an average concentration of 1.53 ± 0.71 ppm. 10 6.1.3. Adverse effects of H2S(g) As mentioned before, H2S has a very typical smell of rotten eggs and can be smelled by the human nose at concentrations as low as 0.5 ppb (Busca and Pistarino, 2003; Firer et al., 2008; Ramírez-Sáenza et al., 2009). It is the most characteristic bad odor constituent in biogas and in the vicinity of anaerobic digesters and wastewater treatment facilities in general (Noyola et al., 2006). At concentrations higher than 100 ppm, the olfactory system is affected and the human nose cannot sense the typical smell (Busca and Pistarino, 2003; Firer et al., 2008). Additionally, hydrogen sulfide is also highly toxic and may be lethal to mammals. When inhaled, H2S(g) inhibits an enzyme which has an important role in mitochondrial respiration (Firer et al., 2008). Exposure of human beings to low H2S concentrations can cause headaches, nausea and irritation of eyes and throat as well as rhinitis, Keratoconjuntivitis, photophobia, intense cough and bronchopneumonia. High concentrations can cause paralysis of the breathing system, unconsciousness and finally death (Merck, 1996 quoted by Noyola et al., 2006). Lethal doses, depending upon exposure, can range from 300 to 1000 ppm. Clinical cases of sulfide poisoning typically involve levels from 4 to 4,200 ppm. In Table 3 some health effects of H2S, with respect to exposure concentrations, are presented. It should be mentioned that H2S is highly toxic not only to mammals but also to aquatic species and vegetation. Apart of being a nuisance and a health problem, the presence of H2S(g) in gaseous fuels may cause operational difficulties. Combustion of a gaseous fuel that contains H2S may lead to the formation of acid rain, due to oxidation of H2S(g) in combustion to sulfur dioxide (SO2). Through heterogeneous oxidation on particulate matter, SO2 is further oxidized to SO3. The SO3 is highly soluble in water that results in the formation of acid rain (Malhotra et al., 2002; Pandey et al., 2004). H2S is a corrosive compound that attacks different materials - iron, copper, cement etc. (Noyola et al., 2006). Thus, fuel gas with high content of H2S expedites corrosion of engines, pipelines and biogas storage structures (Koros and Mahajan, 2000; Chung et al., 2006; Fortunya et al., 2008) Consequently H2S(g) has to be removed from gaseous fuels prior to combustion. 11 Table 3. Hydrogen Sulfide: established dose-effect relationships (Source: WHO, 2000) H2S concentration mg·m-3 ppm 1400 - 2800 1000 – 2000 effect Immediate collapse with paralysis of respiration Reference WHO, 2000a Strong CNS stimulation, 750 – 1400 530 - 1000 hyperpnoea followed by WHO, 2000a respiratory arrest 450 – 750 320 – 530 210 – 350 150 – 250 70 – 140 50 – 100 15 - 30 10 - 20 Pulmonary oedema with risk of WHO, 2000a; Firer et al., 2008 death Loss of olfactory sense Serious eye damage Threshold for eye irritation WHO, 2000b WHO, 2000b; Firer et al., 2008 WHO, 2000b a – adopted from: Hydrogen sulfide. Geneva, World Health Organization, 1981 (Environmental Health Criteria, No. 19). b – adopted from: Savolainen, H. Nordic expert group for TLV evaluation. 40. Hydrogen sulfide. Arbeta och hdlsa, 31: 1-27 (1982). 6.1.4. H2S(g) threshold concentration The World Health Organization (WHO) recommends a H2S(g) guideline value of 2.1 ppm (0.15 mg/m3) for an average exposure time of 24 hours, to ensure no eye irritation. Another recommendation of the WHO is a threshold of 10 ppb (7 µg/m3) for a 30-minute average exposure period, to avoid odor annoyance among the exposed populations (WHO, 2000). The US Occupational Safety and Health Administration (OSHA) established an acceptable ceiling concentration of 20 ppm at the workplace, with a maximum concentration of 50 ppm for no more than 10 min. The US National Institute of Occupational Safety and Health (NIOSH), and the Japanese and Swedish equivalents, have set a maximum recommended exposure limit ceiling value (10 min) of 10 ppm (Lawrence et al., 2000; Busca and Pistarino., 2003; Ramírez-Sáenza et al., 2009). 12 Considering the progressive nature of legislation and the pressure exerted by the legislators as a result of public awareness, it is anticipated that increasingly stricter regulations will be applied in the near future (Iliuta et al., 2004). At present, no common standard has been defined for biogas upgrading to natural gas, but it can be assumed that the European standards will be de-facto the guidelines for most countries. Nevertheless, H2S concentration in biogas should be kept below 7 ppm (5 mg/m3) (Noyola et al., 2006). Normally the H2S specification for the product gas will vary between 4 and 500 ppm H2S, depending upon further use (ter Maat et al., 2005). A concentration of 4 ppm of H2S in natural gas must be reached before the gas is put into a pipeline (DeBerry, 1997). From all of the above it is clear that the removal of H2S from gaseous streams is required for the health of the general public, occupation safety and operational reasons. While concentrations as high as 50,000 ppm exist in gaseous fuels, threshold concentrations are around 10 ppm. 6.2. Methods for H2S(g) removal The removal of hydrogen sulfide from the gaseous streams depends on various factors e.g., raw feed composition, treated gas quality, economic analysis of the process of desulphurization, and the corrosion problems in existing operational units. The gaseous streams, therefore, require desulphurization through a techno-economicallyviable process in order to meet the product purity requirement, to generate clean fuels (in case of gaseous fuels) and also to conform to the stringent sulfur emission standards (Pandey et al., 2003). Most of the purification and treatment processes of gaseous streams that contain H2S(g) result in generation of sulfate (SO42-) or elemental sulfur (S0) through oxidation. Elemental sulfur is a non-corrosive solid that is easy to handle and transport. In addition, it has a commercial value exceeding that of sulfuric acid, although both are used in chemical processing and fertilizer production (Qaisar et al., 2007). Most of the elemental sulfur is produced to-date in oil refineries and natural 13 gas treating plants, and is sold for the production of sulfur compounds such as sulfuric acid (Busca and Pistarino, 2003). Technologies for gaseous pollution control are varied. Fig. 1 shows the common application of each type of technology based on pollutant concentration and air or gas flow. Many of the H2S(g) removal processes are based on its absorption into an aqueous solution followed by further treatment. Other processes are preformed in the gaseous phase or involve a solid phase, as described in the following paragraph. Figure 1. Applicability of various gaseous pollution control technologies based on gas flow rates and concentrations to be treated (Source: Noyola et al., 2006). 6.2.1. Gaseous phase and solid phase processes The most common gaseous phase process is the Claus process. It allows the recovery of sulfur from highly concentrated streams, and is reported to be reliable at a sulfur production rate higher than 20 tons per day (Busca and Pistarino, 2003). The process is based on the oxidation of hydrogen sulfide to elemental sulfur, as described in Eq. (3): H 2 S + 12 O2 → S 0 + H 2O (3) The process consists of two steps (Eq. (4) and (5)): step I: H 2 S + 32 O2 → SO2 + H 2 O (4) step II: 2 H 2 S + SO2 → 3S 0 + 2 H 2 O (5) 14 The typical Claus process industrial configuration is comprised of two steps. In the first step one-third of the H2S concentration is first oxidized to SO2 in a burner. A second exothermic step is performed at lower temperatures, namely 473–573 K, in a series of fixed bed catalytic reactors in the presence of alumina as a catalyst. However, one step processes can also be performed. The produced sulfur is condensed and stored (Busca and Pistarino, 2003). Numerous desulphurization plants based on Claus and Super Claus processes are in operation worldwide (Pandey et al., 2004). However, the need for high temperature makes this process economically feasible mainly for systems that work already at high temperatures. Other gaseous phase processes are thermal incineration and catalytic combustion, which are used to remove highly concentrated VOC-containing streams, with a benefit of additional energy supply. H2S(g) in the gas stream, if exist, will be oxidized into sulfur dioxide (SO2), that has a much less pungent irritating odor with a much higher recognition threshold (three orders of magnitude higher) than sulfide compounds. Nevertheless, sulfur dioxide is also a toxic substance and its emissions are regulated. The temperatures needed for these processes are high: above 1000K for thermal incineration and 600-800K for catalytic combustion (Busca and Pistarino., 2003). Adsorption on solids is a common technology to recover volatile compounds from contaminated air or gases. The pollutant molecules contact the surface of a solid adsorbent and bond via weak intermolecular forces. The most used adsorbents for H2S(g) removal are activated carbon and zeolite, as well as iron oxide based materials. Adsorbent regeneration is accomplished by volatilization of the adsorbed compounds, stripping (commonly with steam or nitrogen) or by purging with a solvent (Busca and Pistarino, 2003; Noyola et al., 2006; Cosoli et al., 2008). The outlet stream from the regeneration step contains much higher concentration of H2S than in the inlet stream, which still needs to be annihilated. 6.2.2. Physical-chemical processes in the aqueous phase Aqueous phase processes are usually preformed in near-ambient temperature, which means that they require less energy than gas phase and solid phase processes. 15 Most of the H2S(g) removal methods in the aqueous phase are based on Reactive absorption, i.e. absorption of the gas in an aqueous solution accompanied by a chemical reaction. Absorption is usually done through scrubbing. In a scrubber, transfer of pollutants from a gas stream to an aqueous phase is accomplished by intense contact of the polluted gas with water or an absorbent solution, within a packed column or a spray tower. Mass transfer depends on the concentration, the air/water partitioning (Henry law) coefficient and the mass transfer resistance of the scrubber system (Noyola et al., 2006). Scrubbing is reasonably applicable to gases containing a high concentration of H2S to recover elemental sulfur through oxidation, and is largely applied in refineries, for example, to recover H2S from gases arising from hydrodesulphurization processes (Busca and Pistarino, 2003). Different absorbing solution can be used, as presented hereunder. Caustic scrubbing - Absorption is favored by highly alkaline conditions as can be concluded from the dissociation constants of H2S(aq) (Eq. (1) and (2)). A gas stream containing the pollutant is fed to an absorption tower with high alkalinity (i.e., NaOH 50% by weight, pH >12). The absorbent is not regenerated in this process, which requires high reagent consumption and a proper final disposal of the spent solution (Noyola et al., 2006). Chemical precipitation with FeCl2 - The H2S contained in the gas stream is absorbed in a scrubber with a solution of FeCl2 and the dissolved H2S is precipitated as FeS according to the following reaction: Fe 2+ + H 2 S → FeS ( s ) + 2 H + (6) The Fe+2 is not reused in this process, which means considerable reagent consumption (Noyola et al., 2006). Chlorine oxidation - After H2S has been absorbed in a scrubbing tower, it may be oxidized with sodium hypochlorite to produce either elemental sulfur or sulfate, depending on pH, according to the following reactions: HS − + OCl − → S 0 + OH − + Cl − − HS + 4OCl − → SO 4 2− at pH < 7.5 + H + + 4Cl − at pH > 7.5 (7) (8) Chlorine is not reused in the process, so it may result in a high operational cost. Moreover, in the presence of organic compounds, chlorine oxidation is not an 16 attractive oxidation agent because of extended chlorine demand and the formation of undesirable organic chloride compounds (Noyola et al., 2006). Ozone oxidation - H2S or VOCs are dissolved in water within a scrubbing tower and then they are oxidized by ozone. The H2S oxidation is practically instantaneous. Since ozone is instable, its generation should be in situ, which may become an obstacle. In addition, ozone is expensive (Noyola et al., 2006). HS − + O3 → S 0 + OH − + O2 − HS + 4O3 → SO4 2− (9) + H + + 4O2 (10) Potassium permanganate oxidation - After H2S is scrubbed, it can be oxidized using potassium permanganate. This method is not attractive since it has a high cost and the manganese oxide must be adequately disposed of to avoid a negative environmental impact. Different sulfur compounds are produced depending on pH (Noyola et al., 2006). 3H 2 S + 2 KMnO4 → 3S 0 + 2 H 2 O + 2 MnO2 + 2 KOH at pH < 7.5 (11) 3H 2 S + 8KMnO4 → 3K 2 SO4 + 2 H 2 O + 8MnO2 + 2 KOH at pH > 7.5 (12) Hydrogen peroxide oxidation - The oxidation rate of sulfide with hydrogen peroxide is relatively slow: 20 to 30 minutes contact time is normally required for a complete reaction. The mechanisms of oxidation of H2S by hydrogen peroxide are not well understood; however, it is suggested that direct oxidation of sulfide by hydrogen peroxide depends on the reaction with oxygen released during gradual decomposition of hydrogen peroxide (Noyola et al., 2006). H 2 O2 + H 2 S → S 0 + 2 H 2 O at pH < 8.5 2 H 2 O2 + S 2− → SO2 2− + 2H 2 O at pH > 8.5 (13) (14) All the removal processes described above may be highly effective, but all have a common and significant drawback – a high reagent consumption, which usually means high operational costs and the need for a proper disposal way for the waste. Another promising group of techniques for hydrogen sulfide removal are the processes that include regeneration of the reactive agent. 17 6.2.2.1. Regenerative gas scrubbing In case of regenerative processes the reagent used to capture H2S can be recovered. Different solvents that have a high affinity for H2S can be applied, such as ethanolamines seem (Busca and Pistarino, 2003; Fortunya et al., 2008). An example of this process is described by the following reaction (Eq. (15)): R 2 NH + H 2 S ↔ R2 NH 2 HS + heat (15) The heat produced by the exothermic reaction during the absorption step is used to preheat the absorbent in the desorption step, since desorption is favored at higher temperatures. By this means, the absorbent can be regenerated. However, H2S as well as CO2 are discharged as exhaust product, so additional treatment processes would be needed before final disposal (Noyola et al., 2006). Although these processes have been extensively and successfully applied, they have many drawbacks such as high energy and operating costs due to the regeneration of the absorbent phase (Fortunya et al., 2008). Additionally, treatment with solvents like ethanol-amines leaves concentration of residual H2S in the gas in the order of few ppm. This means that deodorization is not achieved without further dilution or treatment (Busca and Pistarino, 2003). Another regenerative method is based on the precipitation of H2S(g) with metal ions accompanied by subsequent regeneration of the solid metal sulfide formed, as suggested by ter Maat et al. (2005). As mentioned earlier, the dissolved H2S molecule acts as a diprotic acid. Sulfides of most bivalent metal ions, e.g. zinc, copper, silver, lead, magnesia, nickel and tin are highly insoluble. Therefore aqueous solutions that contain these metal ions can be used as a washing liquid in desulphurization processes. The solid produced is separated from the spent absorbent, oxidized (into CuO, for example) and then dissolved back in the spent absorption liquid. Another, very common, regenerative scrubbing process is the LRSR process. 6.2.2.2. LRSR process Liquid redox sulfur recovery (LRSR) processes offer inherently good hydrogen sulfide removal and highly flexible operating properties at near-ambient temperatures (nominally 20-50°C). The H2S is converted to elemental sulfur, that has a commercial value, and the consumption of chemicals is minimal since the process is regenerative. 18 The overall reaction in LRSR processes is the same as for Claus process H2S conversion (Eq. (3)) (DeBerry, 1997). However, direct oxidation of H2S by oxygen at ambient temperatures is slow and side reactions tend to form undesirable sulfur oxyanions as byproducts. Therefore, oxidation of H2S in the LRSR absorber is carried out by an intermediate redox couple (“catalyst”) denoted Ox (for the oxidized form of the redox couple) as described in Eq. (16). 2Ox + H 2 S ( aq ) → S 0 + 2 R − + 2 H + (16) The symbol R- is used for the reduced form of the redox couple, which is a one electron redox couple in this example. The reduced form of the redox couple produced in the reaction above is regenerated with air in the oxidizer, according to the overall reaction: 4 R − + O2 ( g ) + 4 H + → 4Ox + 2 H 2 O (17) As shown, this reaction also consumes hydrogen ions, neutralizing the hydrogen ions generated in the sorption step (Eq. (16)). The sulfur slurry is removed from the oxidizer and sent to a filtration unit for removal as a wet (approximately 50% solids) cake (DeBerry, 1997). As the catalyst for the LRSR process the following redox couples can be used: V5+/V4+; Fe3+/Fe2+; Co3+/Co2+ and As5+/As3+. Due to high oxidation rates of H2S by Fe3+ (DeBerry, 1997; Hua et al., 2001) and the fact that the cations of vanadium (V), cobalt (Co) and arsenic (As) are toxic and dangerous to the environment, the iron couple appears to be the most suitable for the LRSR process. The use of vanadium (the so-called Stretford process) has dominated in the 1980's, but stopped spreading because of the chemical degradation and environmental problems associated with vanadium. The iron based processes is gradually replacing the vanadium systems (Hua et al., 2001). More than fifty plants based on Vanadium catalyst are in operation for desulphurization of gaseous streams containing H2S. Similarly, around forty plants all over the world based on iron based liquid-redox processes are in operation (Pandey et al., 2004). The continuing search for more economical methods has led to investigations into microbiological solutions for purifying H2S-containing gases, as well as for the desulfurization of coal and petroleum. Biological methods for H2S(g) removal are discussed in the next section. 19 6.2.3. Biological processes in the Aqueous phase Biological processes are characterized by low capital costs and low energy requirements (Malhotra et al., 2002; Pandey et al., 2003; Ramírez-Sáenza et al., 2009), since they can proceed around ambient temperatures and at atmospheric pressure (Qaisar et al., 2007). A variety of biochemical processes using various bacterial species are reported in the literature that are capable of removing sulfur in various forms from gaseous emissions. In these processes, the microorganisms metabolize the H2S, as a source of energy for growth and maintenance, producing sulfate or elemental sulfur (Malhotra et al., 2002; Noyola et al., 2006). In order to make a good selection of a treatment method, flow rate, type of pollutant and its concentration must be considered. Bio-filters, bio-scrubbers and bio-trickling filters have been proven to be a suitable, environmentally friendly and cost-effective alternative for waste gas treatment, especially for the treatment of low concentrations of H2S (Fortunya et al., 2008). 6.2.3.1. Biological methods Biofilters – The contaminated gas is continuously fed into the biofilter, while a nutrient solution is discontinuously added. The filter bed contains microorganisms growing as a biofilm on the surface and crevices of the support. The H2S(g) is absorbed from the gas to the aqueous bio-film and subsequently oxidized by appropriate bacteria populations. The support media should have high porosity, high buffer capacity, high nutrient availability, and especially high moisture retention capacity (Qaisar et al., 2007). This technique is becoming widely accepted due to the high processing efficiency at low sulfur concentrations, the moderate capital costs, and the very low maintenance costs. Mixed micro-organism cultures naturally grow on appropriate natural biofilter beds so abatement of many volatile compounds can be obtained simultaneously (Busca and Pistarino., 2003). Biotrickling filters - The operational principle of a biotrickling filter is similar to that of a biofilter. In this device, polluted air is passed through a packed non-submerged column where liquid is continuously down-flow recirculated through the packing. The pollutant is first absorbed in the falling liquid film and transferred to the microorganisms that grow attached to the surface of these supports. The liquid provides moisture, nutrients, pH control to the biofilm and allows the removal of 20 inhibiting products. Eventually, excess biomass is sloughed off by the trickling liquid and stable operation can be achieved (Noyola et al., 2006). Bioscrubbers - Firstly H2S(g) from the gas phase is absorbed in a recycling water stream followed by biological oxidation of H2S in the liquid. Nutrient addition and pH are continually controlled in the bioreactor in order to maintain microbial growth and high activity. The excess biomass and byproducts are purged from the system (Noyola et al., 2006; Qaisar et al., 2007). Membrane bioreactors - In a membrane bioreactor, the H2S(g) is transferred through a membrane to a biofilm attached on the other side of the barrier, where nutrients and oxygen are provided. A distinct characteristic of membrane bioreactors is the fact that the polluted gaseous stream and the biomass is physically segregated which allows the use of waste gas treatment in certain extreme applications such as indoor air (Noyola et al., 2006). 6.2.3.2. Microorganisms of particular interest for H2S(g) removal The use of microorganisms to oxidize H2S, producing sulfate or elemental sulfur as a consequence of complete or incomplete metabolism, respectively, has been considered a potential alternative for application on a large scale (Zhang and Tong, 2006). Among the H2S oxidizing microorganisms, Thiobacillus seems to be particularly suited for engineering applications due to its simple nutritious requirements, its high effectiveness and resistance to toxic substances and the wide pH range it can tolerate (Noyola et al., 2006). The most common reaction is a direct oxidation of sulfide to sulfur and sulfates by means of oxygen provided by air (obligate autotrophs). In other cases (Thiobacillus denitrificans) nitrate reduction to N2 allows the oxidation of sulfide to sulfate. Particularly, Acidithiobaillus ferroxidans raises a very simple and effective process for H2S treatment in which the oxidant is regenerated by the microorganisms (as part of the Fe3+/Fe2+ LRSR process) (Noyola et al., 2006). Certain photosynthetic bacteria belonging to families Chromatiaceae and Chlorobiaceae are also being used to metabolize H2S effectively (Malhotra et al., 2002; Mesa et al., 2004). The conventional physicochemical methods for removing hydrogen sulfide from gaseous streams require large investment and operational costs e.g. high pressures, high temperatures or special chemicals. Due to this high costs pre-treatment of 21 gaseous fuels contributes significantly to the overall operation and maintenance costs of any energy recovery system. In addition, secondary hazardous wastes are generated in most processes (poisoned catalysts, contaminated reactor liquids and corroded reaction vessels), which need to be treated. (Malhotra et al., 2002; Pandey et al., 2004; Qaisar et al., 2007). Microbiological processes, on the other hand, tend to be efficient only at low H2S(g) concentrations. Additionally, biological systems need to be fed continuously with nutrients and humidity (Mesa et al., 2004; Chung et al., 2006). If the wanted microbiological population is washed out for some reason (e.g., blowdown of the system, strengthening of another population) it takes time to redevelop the population. From all the H2S(g) removal methods reviewed above, the Fe2+/Fe3+ LRSR process seems to be more cost effective than the other physical-chemical processes and more reliable than the common biological processes. The current research focuses on the Fe3+/Fe2+ LRSR process as an efficient method for H2S(g) removal. Note that from this point onward the total-iron species in the oxidation state of 3+ and 2+ will be addressed as ferric (Fe(III)) and ferrous (Fe(II)) respectively. The signs Fe3+ and Fe2+ will represents the non complexed ions only. 6.2.4. The Fe(III)/Fe(II) LRSR process As mentioned earlier, the iron couple is advantageous for the LRSR process because of its high reaction rate and low toxicity (compared with vanadium, cobalt and arsenic couples). The H2S(aq) is oxidized by Fe(III) to elemental sulfur according to the following equation: 2 Fe( III ) + H 2 S ( aq ) → S ( s ) + 2 Fe( II ) + 2 H + (18) The regeneration of Fe(III) is basically done by oxidation of the Fe(II) by O2: 2 Fe( II ) + 12 O2 + 2 H + → 2 Fe( III ) + H 2 O (19) In order to achieve an efficient and continuous process, the flux of ferrous should be similar to the flux of ferric. The flux rates depend on the generation rates of Fe(II) and Fe(III) in the process. Solution pH is often a key variable in LRSR processes, since it has a major effect on the generation rates of Fe(II) and Fe(III) and consequently on other parameters of the process. 22 6.2.4.1. Execution of LRSR processes at neutral to moderately high pH levels Hydrogen sulfide is the acidic species of the sulfide weak-acid divalent system (Eq. (1) and (2)), thus its absorption is preferable in alkaline solutions. In addition, spontaneous oxidation of Fe(II) is very slow below pH 6 (Stumm and Morgan, 1996). The downside is that at pH levels above pH3 Fe(III) tends to rapidly precipitate as iron oxides or iron oxy-hydroxides. The use of organic chelating agents, which have the capacity to bond with cationic ions, proved to be efficient to prevent iron precipitation over a wide pH range, typically up to 10, and to afford satisfying regeneration rate of Fe(III) (Iliuta et al., 2004). The main chelating agents in commercial use are amino and polyamino-polyacetic acids e.g. nitrilo-triacetic acid (NTA), ethylene dinitrilo-tetracetic acid (EDTA) and N-(2-hydoxyethyl) ethylenediamine-N,N',N' triacetic acid (HEDTA) (Hua et al., 2001). DeBerry (1997) concluded that iron-chelate processes are probably best operated above pH7 from the standpoint of the initial Fe-S complex formation and sulfur formation. Other factors to be considered in an LRSR process with chelating agents are detailed below. H2S reactive-absorption efficiency - For good absorption efficiency and small scrubber size the rate of reaction of the redox catalyst with H2S should be as high as possible. The more successful catalysts appear to rapidly form a complex with H2S and then more slowly convert it to sulfur. The initial complex formation reduces the equilibrium backpressure of H2S. The initial reaction rates of H2S with several oxidizing agents are shown in Table 4. Table 4. Rates of some oxidizing agents with Hydrogen sulfide (Source: DeBerry, 1997) Oxidant O2 H2O2 Second order rate constant k (M-1sec-1) Comments / source 1.5 × 10-3 For HS-+O2HS+O2-; 20ºC; Resch et al., 1989 -1 Resch et al., 1989 -1 4.5 × 10 V(V) 4.2 × 10 20ºC; Radian FeOH2+ 1.5 × 103 20ºC; pH2; Asai et al., 1990 Fe(III)EDTA 6.2 × 103 20ºC; Radian 23 Also important is the capacity of the scrubbing liquid for H2S. Use of a concentrated redox catalyst can increase the capacity of the sorbent for H2S. However, high catalyst concentrations could also lead to fast precipitation of sulfur in the scrubber and consequent plugging (DeBerry, 1997). Control of sulfur formation is a key topic, as discussed hereunder. Sulfur Formation - The formation of solid sulfur is a key step in LRSR processes. The efficient formation of a separable solid product is important to economical continuous operation of these processes. However, premature formation of sulfur in the H2S scrubber can cause plugging of the scrubber internals (internals are necessary to get good gas/liquid mass transfer and high removal efficiency). Faster oxidation of H2S has been reported to lead to faster precipitation of S0, and hence to smaller particle size of the S0 precipitates and settling difficulties (DeBerry, 1997). Redox Catalyst Regeneration - The reduced form of the redox catalyst must be converted back to the oxidized form in the regenerator, which is usually an airsparged vessel. This vessel is usually large compared to the other vessels used in LRSR processes. In addition, the air compressors and energy needed to run the compressors can be significant. Thus, rapid regeneration reaction rates are needed for economical operation. Generally speaking, the higher oxidation potential the oxidizing catalyst has, the slower and more difficult will be the regeneration step. This is because as the redox potentials get more positive, they approach the redox potential of oxygen, resulting in a decreased driving force for the regeneration of the redox catalyst by O2. Thus, the catalyst redox potential should be high enough to give efficient H2S scrubbing (as discussed above) but low enough to allow efficient regeneration by air. The chelate molecule plays a major role in the rate of regeneration of Fe(III) from Fe(II) in the chelated iron LRSR processes. The air oxidation rate of uncomplexed ferrous ion (Fe2+) is many orders of magnitude slower than that of Fe2+ chelated with EDTA or similar chelating agents. However, Fe(III) chelated with EDTA is a much weaker oxidizing agent than simple aquated (complexed only by water molecules) Fe(III). Oxidation rates of several compounds by oxygen are compared to the rate of oxidation of Fe(II)EDTA in Table 5. The second order rate constants were derived primarily from the overall or the slowest step. The rate constants in Table 5 range over six orders of magnitude. Note that both the uncatalyzed oxidation rate of H2S by O2 and the rate of uncomplexed Fe(II) oxidation by O2 with are very low (DeBerry, 1997). 24 Table 5. Comparison of rates of oxidation of species by Oxygen (Source: DeBerry, 1997) Reactant H2 S Second order rate constant k [M-1sec-1] Comments / Adopted from 1.5 × 10-3 Resch, 1989 pH 8.8; Radian* V(IV) 1.0 Fe2+(aq) 4.0 × 10-4 Lacey, 1970 Fe(II)EDTA 2 2.8 × 10 Radian* Fe(II)DTPA 1.4 × 101 Radian* Fe(II)HEDTA 6.0 × 102 Radian* * Radian Corporation, Austin, Texas Consumption of Chemicals - There is several sources of chemical consumption in liquid redox processes. One of these, carry-out of chemical reagents with the sulfur product, is in large part a “physical” process, although some chemical binding to the surface of the sulfur may occur. A major problem for iron chelate-based processes is chemical degradation of the chelating agent. The degradation is generally thought to be caused by free OH● and H2O2 radicals or other highly reactive intermediates generated in the air oxidation reactions. A number of compounds have been tested in an attempt to scavenge the hydroxyl radicals and thus protect the chelating agents. Thiosulfate (S2O32-) is known as an effective scavenger for the hydroxyl radicals and has been proposed as a prospective degradation inhibitor. However, most of these scavengers are consumed more or less rapidly and thereby lead to a buildup of soluble salts in solution. This will eventually necessitate a blow-down or a solution cleanup step. Eventually, the addition and consumption of chelate agents and radical scavengers elevate the cost of the LRSR process (DeBerry, 1997; Hua et al., 2001). Another source of loss is blow-down of solution necessitated by byproduct formation (see below). Byproduct Formation Rate - Oxidation of H2S beyond the oxidation state of elemental sulfur results in the formation of soluble sulfur oxyanions which are difficult to purge from the LRSR process stream. If these salts reach high concentrations, they can cause scale formation and operating problems. The salts must be removed by “blowdown” or chemical treatment of the solution which results in loss of some of the catalyst and economic penalties. One of the main sources of this problem is the slow 25 and/or incomplete oxidation of sulfide, which leads to elevated concentrations of polysulfides in solution. These polysulfides are readily converted to sulfur oxyanions when the solution contacts air in the regenerator. Chelated iron processes, with respect to vanadium/ADA processes, have much faster sulfide-to-sulfur conversion kinetics, and the concentration of polysulfide in these solutions is usually quite low. Paradoxically, one of the soluble sulfur oxyanion byproducts, thiosulfate, is a desirable constituent of current liquid redox processes since it is a very good inhibitor of chelate degradation, as described above. Thus the control and management of thiosulfate formation is a part of the operating protocol of such processes (DeBerry, 1997). To summarize, the use of organic chelating agents allows faster kinetics for conversion of H2S to sulfur and faster regeneration of the catalyst with oxygen with respect to earlier liquid redox processes. One unfortunate byproduct of this evolutionary step is catalyst degradation, which is controlled in large part by inhibitors and management of solution chemistry. However, most of these inhibitors are consumed rapidly. Another byproduct of the faster kinetics associated with the new catalysts is difficulty in controlling the formation of solid sulfur. In order to avoid the use of chelates and radical scavengers it is possible to perform the LRSR process with the iron couple at low pH levels. This is discussed in the next paragraph. 6.2.4.2. Execution of LRSR processes at low pH levels At low pH levels (pH1 to pH2) the solubility of ferric species is high enough to avoid the need for chelating agents. Since the rate of ferrous oxidation is very low at this pH levels (see Subsection 1.2.4.1), bacteria are added to catalyze the regeneration process. The most commonly used bacteria are Acidithiobacillus Ferrooxidans (A.F.). These are a gram-negative rod-shaped, chemo-autotrophic, mesophilic and acidophilic bacteria, with the ability to oxidize Fe(II) in acidic solutions and couple the energy to support carbon dioxide fixation and growth. This ability is particularly suited for regeneration of Fe(III) in the LRSR process. The A.F. bacteria grow at pH values between pH 1 and 4.5, and the optimal pH and temperature ranges for ferrous iron oxidation are 1.8–2.5 and 30 –35 °C, respectively (Malhotra et al., 2002; Mesa et al., 2004; Chung et al., 2006). In 26 Addition, the problems associated with other H2S oxidation microbiological processes are avoided as H2S does not have an inhibiting effect on A.F. and SO42- is not accumulated in the medium (Noyola et al., 2006 ; Ramírez-Sáenza et al., 2009). The biological LRSR process at low pH levels consists of two steps, usually done in two different reactors. The first one is a scrubbing reactor, where H2S(g) is absorbed into ferric sulfate (Fe2(SO4)3) solution and oxidized to S0, while Fe(III) is reduced to Fe(II) (Eq. (20)). The solution from the first reactor flows to the second reactor, an aerobic bioreactor, where Fe(III) is regenerated from Fe(II) by A.F. bacteria (Eq. (21)). The ferric iron is then recycled into the first reactor to repeat the cycle (Malhotra et al., 2002; Ebrahimi et al., 2003; Mesa et al., 2004; Pandey et al., 2004; Chung et al., 2006). H 2 S + Fe2 ( SO4 ) 3 → S 0 + 2 FeSO4 + H 2 SO4 (20) 2 FeSO4 + H 2 SO4 + 0.5O2 → Fe2 ( SO4 ) 3 + H 2 O (21) The bioreactor can be operated either as a fixed or fluidized bed. The process is schematically depicted in Fig. 2. Figure 2. Process scheme of a chemo-biological process for H2S(g) removal (Source: Ebrahimi et al., 2003) A distinct advantage of the process is that the reaction of H2S with Fe2(SO4)3 is so rapid and complete that it does not produce toxic waste or generate other waste products apart from sulfur (Ebrahimi et al., 2003; Mesa et al., 2004). The overall process rate is of a similar order of magnitude as that of the alkaline LRSR process with chelating agents (Hua et al., 2001). Hence, the operating cost of the chemobiochemical processes is extremely low as compared with other chemical redox processes for desulphurization of the gaseous fuels (Pandey et al., 2003). In general, 27 the operation costs of this process are around one third of those of conventional processes such as scrubbers and absorption columns (Noyola et al., 2006). In order to achieve an efficient and continuous process, the flux of Fe(II) from the absorbent reactor should be similar to the flux of Fe(III) from the bioreactor. The Fe(II) generation step in the absorption reactor depends on two reactions occurring one after the other: H2S(g) absorption and H2S(aq) oxidation (i.e., reactive absorption). Thus the Fe(II) generation is related to the efficiency of both reactions. Consequently, the main factors that affect Fe(II) generation, as reviewed in literature are: the H2S(g) concentration, the Fe(III) concentration and pH. These factors are all reported to have a positive effect on the Fe(II) generation rate. The positive effect of H2S(g) flux on Fe(II) generation was reported by Asai et al. (1990), Ebrahimi et al.(2003), Pandey et al. (2003), Chung et al. (2006) and Gendel (2007). Most of these works found a linear relationship between H2S flux and ferrous generation rate, as shown in Fig. 3. Figure 3. Relationship between inlet H2S(g) loading and Fe(II) production rate (Source: Chung et al., 2006). The positive effect of Fe(III) was reported by Asai et al. (1990), Ebrahimi et al.(2003), Pandey et al. (2003), Mesa et al. (2004), Pandey et al. (2004), Chung et al. (2006) and Gendel (2007). Asai et al. (1990) reported that the species which reacts with H2S is FeOH2+ and that the reaction is irreversible and first order in both H2S and FeOH2+. The overall reaction is: H 2 S + 2 FeOH 2 + → S 0 + 2 Fe 2 + + 2 H 2 O 28 (22) which consist of the following steps: step I: H 2 S + FeOH 2+ → H 2 S • FeOH 2+ (23) step II: H 2 S • FeOH 2+ + FeOH 2+ → S 0 + 2 Fe 2+ + 2 H 2 O (24) Asai et al. (1990) assumed that FeOH2+ was the sole reactive species in the system. However, Ebrahimi et al. (2003) have evaluated the concentration of all possible ferric species by taking into account all of the ferric complex equilibrium equations and showed that other ferric hydroxide species (Fe3(OH)45+ and Fe2(OH)24+) were on the same order of magnitude as FeOH2+ concentration under the experimental conditions (Fig. 4). Thus Ebrahimi et al. (2003) concluded that the reactive-absorption rate of H2S(g) was first order with respect to both H2S(g) and the total ferric iron concentration. Figure 4. Concentrations of different ferric species as a function of pH at 45ºC and 0.2 M Fe2(SO4)3 (Source: Ebrahimi et al., 2003) Pandey et al. (2003) found that further increase in the H2S(g) loading rate resulted in the decrease of its removal efficiency (and consequentially Fe(II) generation) due to limited availability of Fe(III) in the reaction medium. It is also possible that at high 29 Fe(III) concentration the H2S(g) concentration becomes the limiting factor. This may explain the decrease in Fe(II) generation with further increase in Fe(III) concentration, as reported by Ebrahimi et al.(2003) , Pandey et al. (2003), Mesa et al. (2004) and Pandey et al. 2004). Another explanation is that above a certain concentration of Fe2(SO4)3 the ionic strength and viscosity of the solution increase remarkably resulting in a decrease of the solubility, diffusivity and liquid side mass transfer coefficient of H2S and hence absorption rate (Ebrahimi et al., 2003). In the experiments conducted by Ebrahimi et al. (2003) the limiting concentration was 0.3 M. However, this value is different in each work, due to different conditions of experiments. All experiments reviewed were conducted at pH between 1.1 and 2.3. A positive pH effect was reported by Asai et al. (1990), Ebrahimi et al.(2003), Pandey et al. (2003) and Pandey et al. (2004). There are two possible causes: first, with the increases in pH there is an increase in ferric hydroxide species that have a positive effect on Fe(II) generation according to Ebrahimi et al.(2003); second is the poor solubility of H2S in low pH ferric sulfate solutions (Pandey et al., 2003). Two other parameters that affect H2S(g) reactive-absorption and can be modified quite easily are the surface area for gas-liquid mass transfer and the gas retention time in the adsorbing solution. Pandey et al. (2003) investigated the reactive-absorption efficiency of H2S(g) in a packed bed reactor and found efficiencies above 96%. Packing material that had large surface area resulted in high removal efficiency of H2S(g) due to better mass transfer of H2S from the gaseous bulk to the liquid phase. However, the sulfur produced in the process with the packing material with the largest surface area accumulated in the chemical oxidation unit, which resulted in plugging of the reactor. Chung et al. (2006) examined three gas retention times and their results indicated that a longer GRT could elevate H2S(g) removal efficiency. Gendel (2007) reached similar results when he examined two different flow rates of the treated air, and found that at low flow rate the removal efficiency of H2S(g) was higher. As for the Fe(II) generation step in the bioreactor - in all the works that were reviewed an efficiency higher than 90% for Fe(II) oxidation by A.F. was achieved (Malhotra et al., 2002; Pandey et al., 2003; Mesa et al., 2004; Pandey et al., 2004; Chung et al., 2006 and Gendel, 2007). An Fe(II) oxidation rate of at least 11.3 mM/h (0.63 g/l/h) was reported by Mesa et al. ( 2004), Chung et al. (2006) and Gendel (2007). 30 A.F. can convert ferrous ions to ferric ions in the presence of TDS concentration up to 50,000 mg/l and even 60,000 mg/l with proper acclimatization (Malhotra et al., 2002). The presence of elemental sulfur hardly affects the bio-oxidation of Fe(II) to Fe(III) according to Malhotra et al. (2002). When the dissolved oxygen concentration is not a growth limiting substrate i.e., above 0.29 mg/l (Liu et al., 1988, quoted by Gendel, 2007) and for a given pH and temperature, both A.F growth and the ferrous iron oxidation rate depend on three parameters: substrate (i.e. ferrous iron) concentration, product (i.e. ferric iron) concentration (competitive and non-competitive inhibitions) and the biomass concentration. The ratio between the volumes of the chemical (absorption) and biological reactors depends on the ratio between the rates of H2S(g) removal and Fe(II) oxidation. According to Gendel (2007), who examined the removal efficiency of H2S(g) in concentrations between 2 and 57 ppm, the required volume for the biological reactor can be designed three orders of magnitudes smaller than the chemical absorption reactor. However Chung et al. (2006), who experimented with H2S concentrations of 500-1500 ppm, conclude that the volume ratio of biological reactor to chemical reactor should be 13.5:1. All in all, the method of biological regeneration of Fe(III) by Acidithiobacillus Ferrooxidans as part of the LRSR process at low pH suffers from two major drawbacks: (1) the dependence of the process on the sensitive autotrophic biomass and (2) the relatively rapid precipitation of Fe(III) solids at the pH range optimal for A.F. bacteria (mostly as ferric sulfate-hydroxides of the jarosite group). With respect to the first point, it was mentioned earlier (Subsection 1.2.3.2) that biological systems need to be fed continuously and it takes time to redevelop a bacterial population. This is true especially for autotrophic bacteria, as A.F., which are much more sensitive to changes in environmental conditions (e.g. temperature, fluctuation in substrate flux, pH) and have a very low yield coefficient. To exemplify the time needed for the development of A.F. population: in the experiment conducted by Chung et al. (2006) the A.F. bacteria that were used for the experiment were isolated from acid mine drainage, were grown in a growth medium for 7 days and then were cultivated for two further months on the packing material (granular activated carbon). 31 As for the second point, the formation of ferric iron precipitates has been reported in most studies that involve ferrous iron oxidation by A.F. In many of the reports the precipitates were identified as the iron hydroxide jarosite, a basic ferric hydroxylsulfate with the general formula MFe3(SO4)2(OH)6, where M may stand for K+, Na+, NH4+ or H3O+ (Pagella and Faveri, 2000; Mesa et al., 2004). The formation of jarosites depends strongly on pH. Gendel (2007) showed that the mass of precipitate that was formed increased with an increase in the initial pH from 1.25 to 2.5. Although the formation of jarosite was significantly reduced at lower pH values, jarosite precipitation did not cease and remained a major problem for long-term operation even at low pH values. A further decrease in pH (i.e. pH≤1.0), may diminish the problem, however at such pH the oxidation efficiency of Fe(II) by A.F. is very low, and the bacteria require a long acclimation period. The problems associated with jarosite precipitation are the following: (1) The precipitation reduces the soluble ferric concentration in the aqueous phase, and thus less ferric is available for oxidizing H2S in the desulphurization process; (2) Precipitation of ferric compounds tends to block pumps, valves, pipes etc., and causes maintenance difficulties; (3) Precipitation of ferric compunds on the surface area of the carrier in fixed biological beds reduces both the amount of surface area available for biomass growth, and the flux of the substrate (ferrous iron) into the biofilm, thus reducing the overall efficiency of the biological step. In experiment by Mesa et al. (2004) the final precipitate was found to accumulate on the support at rates ranging between 0.1 and 0.8 g/l/h. This implies that a one cubic meter reactor working continuously for 100 days would accumulate approximately 2000 kg of precipitate. Jarosite formation depends also on the ferric iron concentration: when the concentration is below 1000 mg/l jarosite precipitation is considerably slow (Gendel, 2007). However, at low ferric concentration the reactive-absorption efficiency of H2S(g) is low. Only one study was found to describe an LRSR process at low pH, without the use of Acidithiobacillus Ferrooxidans for Fe(III) regeneration. Zhang and Tong (2006) used copper to catalyze Fe(II) oxidation by O2. The experimental conditions were as follows: 0.54 M Fe(II), 1.07 M Fe(III), 1.26 M Cu(II); pH was low but not specified; H2S(g) concentration was around 1000 ppm and the gas (air + H2S) flow was 11.7 32 m3/min (3.0 l/m3 liquid-gas ratio). According to Zhang and Tong (2006) the cupric ion precipitated rapidly as cupric sulfide (Cu(II)S(-II)) (Eq. (25)), however it was then rapidly oxidized by Fe(III) to elemental sulfur (Eq. (26)). The Cu2+ ion has a catalytic effect on Fe(II) oxidation, hence the oxidation of Fe(II) and consequently the regeneration of the absorbing solution were fast. (25) Cu 2+ + S 2− → CuS (26) 2 Fe 3+ + CuS → S 0 + 2 Fe 2+ + Cu 2+ quick quick Zhang and Tong (2006) reported that the H2S(g) removal efficiency of their suggested process was almost 100% during the entire experiment (that lasted 60 minutes) while the concentrations of cupric ion and ferric ion were very close to the initial values. No CuS was found to accumulate. In addition, sulfur selectivity after 4 h of experiment was calculated to be higher than 96% (Sulfur selectivity = Moles of sulfur produced per moles of H2S reacted). This process may be a successful alternative for an LRSR process at low pH, without the disadvantages of a biological system. To sum up all of the above, the methods for H2S(g) removal are varied, and the favored method depends on the gas origin and content. The aqueous phase methods usually have the lowest energy requirements, and among them the Fe(III)/Fe(II)-LRSR method seems favorable since the oxidizing agent is regenerated in the system and the only products are H2O and S0. When conducted at neutral to high pH (7-9) the H2S(g) removal is very fast, but the use of chelating agents is essential in order to minimize precipitation, mainly of ferric hydroxides. These chelating agents tend to degrade with time, a fact which eventually leads to high operation costs. It is possible to perform the Fe(III)/Fe(II)-LRSR method at low pH levels, in order to minimize precipitation, with the use of Fe(II) oxidizing bacteria. In this way the use of chelating agents is avoided, but the process becomes dependent on a very sensitive biomass (autotrophic bacteria). Moreover, some precipitation still occurs, which is critical in continues processes. In this work two alternatives were investigated for Fe(III) regeneration at low pH levels. The first one is chemical oxidation of Fe(II) by O2 in the presence of catalytic reagents and the second is the use of electrolysis to oxidize Fe(II). In both methods the pH would be kept around pH1.0. Thus, the two major drawbacks of the low pH biological LRSR are bypassed: there is no dependence on a biological system and the 33 precipitation problem becomes minor. Before describing the methods in more details the mechanisms of ferrous oxidation by O2 and electrochemical oxidation are extensively reviewed. 6.3. The kinetics of Ferrous iron (Fe(II)) oxidation by oxygen 6.3.1. pH dependence As early as 1906, the rate of oxidation of ferrous ion by dissolved oxygen in near neutral pH solutions was shown to be (Goto et al., 1970): − d [ Fe( II )] = k[ Fe( II )][OH − ] 2 PO2 dt where k = 8.0(± 2.5) × 1013 min −1 atm −1 mol −1liter 2 (27) at 20 0 C . This kinetic equation was confirmed empirically and used in many studies (Sung and Morgan, 1980; Millero, 1985; Stumm and Morgan, 1996 etc.). Stumm and Morgan (1996) compared Eq. (27) with fitted curves obtained from empirical data attained from the oxidation of ferrous solutions with concentrations lower than 0.5 mM/L (~28 mg Fe/l), at the pH range between ~ 4 and 6. In Fig. 5 it can be seen that below pH ~4 the rate of oxidation is very low, and is fundamentally independent of pH. Millero (1985) presented data for the rate of oxidation of Fe(II) in solution which, among other observations, includes measurements at pH values higher than 6 (Fig. 6). The results of Millero (1985) indicate that at pH values greater than ~8 the rate of ferrous oxidation is also independent of pH (these data was obtained with seawater but the author extrapolated it also to freshwater). At pH values lower than ~8, the results of Millero (1985) are in agreement with Fig. 5. Thus the overall Fe(II) oxidation rate versus pH plot shows a central region where the rate is strongly pH dependent, flanked by regions on either side where the rate does not change as a function of pH. Note that in freshwater above pH 9.0 the concentration of soluble Fe(II) species is so low that it is very difficult to obtain reasonably accurate empirical kinetic data (Morgan and Lahav, 2007). 34 Figure 5. Oxidation rate of Fe(II) Figure 6. Rate constants for oxidation of species as a function of pH ( PO2 = 0.20 soluble Fe(II) species in water as a function atm). (Source: Morgan and Lahav) of pH (Source: Millero,1985) The trend apparent from these observations can be explained in terms of the equilibrium chemistry of Fe(II) in aqueous solutions. The following rate equation given by Millero (1985) separates the individual Fe(II) species present in aqueous solution: − d[Fe(II)] − = (k o ([Fe 2+ ] + k 1[Fe(OH) + ] + k 2 [Fe(OH)02 (aq) ] + k 3 [Fe(OH)3 ])DO (28) dt Where: ko, k1, k2 and k3 are oxidation rate constants (time-1). The value of dissolved oxygen concentration (DO) is used instead of PO2 since it is in fact DO which participates in the oxidation reaction and the DO concentration will vary for a given PO2 for a variety of reasons such as temperature, ionic strength and the oxygen consumption rate of the solution (Millero et al., 1987). Eq. (28) implies that in the course of soluble ferrous species oxidation, more than one oxidation reaction occurs simultaneously (Stumm and Morgan 1996), e.g. a Fe 2+ → Fe 3+ + e − b Fe(OH ) → Fe(OH ) + e − c Fe(OH )2 → Fe(OH )2 + e − + 0 2+ + 35 It implies also that the rate of oxidation depends on the concentration of each species and its individual oxidation rate. Fig. 7 is a log species – pH diagram for Fe(II) species in single phase aqueous equilibrium (Total Fe(II) concentration = 10-3 M) and Fig. 8 is a log species – pH diagram for Fe(II) for a two phase, aqueous – solid equilibrium, in the absence of complexing agents for ferrous iron other than water (i.e. in the absence of Cl-, CO32-, SO42-, etc.). Inspection of Fig. 8 reveals that in order to reach saturation at pH values below ~8, the total Fe(II) concentration must be much higher than that used to construct Fig. 5. In other words, below pH ~8 all Fe(II) solutions are effectively nonsaturated solutions and the equilibrium chemistry of Fe(II) solutions below ~pH 8 is therefore better described by Fig. 7 (single phase aqueous equilibrium). For solutions above ~pH 8 Fig. 8 (two phase aqueous-solid equilibrium) can be used to predict the fraction of the ferrous concentration that is in the solid phase (Fe(OH)2(s) in this case) (Morgan and Lahav, 2007). Once formed, the rate of oxidation of Fe(OH)2(s) has been shown to be independent of pH, and to depend only on the rate of the introduction of the oxidant (typically O2) into the water (Prasad and Ramasastry,1974 quoted by Morgan and Lahav, 2007). pH pH 0 2 4 6 8 10 12 0 14 Fe(OH)3- 6 8 10 12 14 Fe2+ log species -2 FeOH+ 0 -9 Fe (O H )2 log species Fe2+ -6 -12 4 0 0 -3 2 Fe(OH)2(s) -4 -8 -15 FeOH+ -6 -10 FeT Fe(OH)20 Fe(OH)3- Figure 7. Log species – pH diagram of Figure 8. Aqueous – solid phase soluble ferrous hydroxide species at equilibrium for soluble ferrous hydroxide infinite dilution (Morgan and Lahav, species at infinite dilution (in the absence 2007) of both inorganic carbon species and ion pairing species such as Cl-, SO42-, PO43etc.) (Morgan and Lahav, 2007) King et al. (1995) developed a detailed model of Fe(II) oxidation in homogenous solution based on the Haber-Weiss mechanism. It is the most widely accepted 36 mechanism to describe the oxidation of Fe(II) by O2(aq) and has been used as the basis for several subsequent studies of iron speciation and redox cycling in natural systems (Santana-Casiano et al., 2006; Pham and Waite, 2008). The mechanism describes four one-electron steps for the reduction of the terminal electron acceptor O2(aq): A Fe( II ) + O2 ( aq ) → Fe( III ) + O2• (−aq ) • B Fe( II ) + O2• (−aq ) + 2 H + → Fe( III ) + H 2 O2 (aq ) C Fe( II ) + H 2 O2 ( aq ) → Fe( III ) + OH (•aq ) + OH − D Fe( II ) + OH (•aq ) → Fe( III ) + OH − = radical Combining sets of equations (a, b, c) with (A, B, C, D), Stumm and Morgan (1996) showed that the free energy changes which occur during the oxidation of Fe(II) by oxygen were more negative for the oxidation of Fe(OH)20(aq) than for either Fe(OH)+ and Fe2+. In all three reactions a, b, c step A was found to be endergonic but was least endergonic for reaction c; and for all three reactions a, b, c steps B, C, D were all found to be exergonic but in each case reaction c was the most exergonic. Thus by explicitly making a kinetic argument from thermodynamic data, Stumm and Morgan (1996) deduced that step A was the slowest step in each case (because it was the most endergonic) and was therefore rate-limiting; and that step A was the fastest for reaction c (because, of the three, it was the least endergonic). Similarly, they concluded that since reaction c is most exergonic for steps B, C, D, all in all Fe(OH)20(aq) produces the fastest oxidation sequence. This conclusion is backed up by two other sources. Firstly, it has been shown that ‘hydrolyzed’ ferrous iron species are more readily oxidized than non-hydrolysed ferrous species in the following order Fe(OH)20(aq) >> Fe(OH)+ >> Fe2+, the reason presumably being that OH- ligands donate electron density through both the σ and π systems to the reduced metal ion which increases reducing power and stabilises the Fe3+ formed during the oxidation (Luther, 1990 quoted by Stumm and Morgan, 1996). This hypothesis was experimentally substantiated (in retrospect) by Millero (1985), who showed that the rate constant for the oxidation of Fe(OH)20 is 5 orders of magnitude higher than the rate constant for FeOH+, which, in turn, is 5 orders of magnitude greater than the rate constant for Fe2+ reported by Lowson (1982). Secondly, Wehrli (1990) (quoted by Morgan and Lahav, 2007) showed a linear free energy relationship between the free 37 energy of the reaction (log K) and the rate (log k) of the reaction for the oxidation of Fe2+, FeOH+ and Fe(OH)20(aq). Based on the studies of Lowson (1982) and Millero (1985) the following explicit rate equation for soluble ferrous hydroxide species oxidation can be written (Eq. (29)). − d[Fe(II)] 6 ⋅ 10- 5 ⋅ Fe(II)T 1.7 ⋅ Fe(II)T 4.3 ⋅ 105 ⋅ Fe(II)T = + + 2 3 + 2 K1K w K1K 2 K w K1K 2 K w (H ) K 2 K w K 2 K 3K w KK (H + ) (H + )2 dt 1+ + + 1+ 3 + w + + + + 2 1 + (H + ) + (H + )2 + (H + )3 K1K w (H ) (H ) (H ) K 2 K w K1K 2 K 2w × DO (29) Where: -dFe(II)/dt is the rate of ferrous oxidation in mole-Fe(II)/l/min and DO corresponds to O2(aq) in equilibrium with PO2 =0.20 bar. Each species is expressed explicitly as a function of the total ferrous iron concentration (Fe(II)T), pH and stability constants. The effect of ionic strength and temperature is determined correspondingly by adjusting the equilibrium constants by the Davis equation (based on the Debye - Huckel approach) and the Van't Hoff equation (Morgan and Lahav, 2007). Note that no rate constant has been reported for the oxidation of Fe(OH)3- probably due to empirical limitations, however the concentration of this species is exceptionally low at pH values below pH ~10 and thus, from a practical standpoint, this rate constant is of a lesser importance. Also note that the rate constants in Eq. (29) should probably be considered only with respect to their order of magnitude, rather than as accurate quantities. Having established the relative rates of oxidation of Fe(II) species one can now rationally explain the oxidation rate curve seen in Fig. 5 and Fig. 6. It has already been established that Fig. 7 applies only to Fe(II) solutions at pH below ~8. Examination of Fig. 7 shows that in this pH range the concentrations of FeOH+ and Fe(OH)20 rise steeply and linearly with pH. Since these species (especially Fe(OH)20) are far more readily oxidized than Fe2+ is, this explains the pH-dependence of the oxidation rate between pH ~4–8. Indeed, above pH ~5.5 both the first two terms on the right hand side of Eq. (28) fall away (Stumm and Morgan 1996), which accounts for the second order dependence of the rate law shown in Eq. (27) for this pH region. Below pH ~4 the concentrations of FeOH+ and Fe(OH)20 are so low as to be negligible and Fe2+ dominates. Since the oxidation of Fe2+ is independent of hydroxyl groups, the oxidation rate is no longer pH dependent. 38 For pH values above ~8, all Fe(II) solutions of the order of at least 10-3 molar concentration (under the assumption that ferrous complexes only with H2O) will be saturated solutions. Fig. 8 shows the speciation of Fe(II) in two phase aqueous – solid equilibrium. Here at pH ~8 although Fe2+ exceeds both FeOH+ and FeOH20 in concentration, the latter two species are far more easily oxidized due to the presence of the hydroxyl groups. As one moves from pH ~8 to the right in Fig. 8, both the Fe2+ and FeOH+ concentrations drop steeply whilst the FeOH20 concentration is constant. As the pH increases above 8 the oxidation rate (Eq. (28)) therefore becomes dominated by Fe(OH)20. Since this species is constant with respect to pH ‘throughout’ the range, the rate above pH ~8 becomes independent of pH. As explained earlier all Fe(II) solutions of concentrations of the order of 10-3 M are unsaturated below pH ~8, thus the Fe(OH)20 concentration is not constant below pH ~8 (as in Fig. 7 and not Fig. 8) (Morgan and Lahav, 2007). 6.3.2. Fe(II) oxidation by O2 in natural waters In the presence of complexing or ion pairing agents other than H2O, as in natural waters, more components should be added to Eq. (28). King (1998) suggested the following oxidation rate: − d [Fe( II )] = [O2 ] × [Fe( II )]× k app dt (30) ( ) where k app = 4 k1α Fe 2 + + k 2α FeOH + + k 3α Fe (OH ) 0 + ...k nα n and αi is the fraction of each 2 Fe(II) species in solution, ki is the second-order rate constant for oxidation by oxygen, and [Fe(II)] is the total or analytical Fe(II) concentration. The factor 4 in Eq. (30) reflects the stoichiometry of Fe(II) oxidation by oxygen when the first oxidation step (according to the Haber-Weiss mechanism) is rate limiting. The oxidation of Fe(II) with O2 has been widely studied for many systems in marine and atmospheric science and hydrometallurgy. Most experiments were conducted with micro-molar and nano-molar concentrations of Fe(II), as found in aquatic systems. Carbonate is an important ligand in all natural waters and may form a variety of complexes with both Fe(II) and Fe(III). 39 King (1998) conducted experiments in pure water with initial Fe(II) concentrations between 1 to 5 micromolar. According to King's (1998) results the ferrous carbonate complexes (FeCO30, Fe(CO3)22-, and Fe(CO3)(OH)-) dominate the speciation of Fe(II) in natural waters containing greater than 1 mM carbonate alkalinity, where FeCO30 is the most dominant species (see Fig. 9). In contrast, the Fe(CO3)22- complex is the most kinetically active species for pH values above 6, while at pH values below 6.0, the oxidation rate of Fe(II) is well described in terms of the Fe2+ and FeOH+ species (Fig. 10). Santana-Casiano et al. (2005) and Pham and Waite (2008) who investigated the Fe(II) oxidation rate at nanomolar levels (0.025 - 0.250 µM) in natural waters, attained the same tendency in their results. Figure 9. Fe(II) speciation in pure water Figure 10. Contribution of specific Fe(II) with 2.3 mm of NaHCO3 and 0.03 atm of species in total Fe(II) oxidation rate by PCO2 . (Source: King, 1998) O2. Calculations are for pure water with 2.3 mM NaHCO3 at 25 °C. (Source: King, 1998) In seawater, chloride and sulphate complexation with Fe2+ and increased ionic strength are the significant processes responsible for reduced rates as compared to pure water (Sung and Morgan, 1980; Lowson, 1982; Millero, 1985; King, 1998). The complexation with Fe2+ decreases the Fe2+ fraction by a factor of 2 and correspondingly shifts all other species by a comparable amount (King, 1998; Santana-Casiano et al., 2006) (see Fig. 11 in comparison with Fig. 9). Since FeCl+ and FeSO4 species have negligible contribution to the oxidation rate it explains the 40 observed rate in seawater, which is usually a factor of 100 times slower compared to solutions of freshwater-like composition (Sung and Morgan, 1980). Figure 11. Fe(II) speciation in 0.7 m of NaCl and 0.03 m of Na2SO4 with 2.3 mm of NaHCO3 and 0.03 atm of PCO2 . (Source: King, 1998) 6.3.3. Fe(II) oxidation by O2 at low pH levels The oxidation rate of Fe(II) by oxygen at low pH levels (pH < 2) has been examined under various conditions. Table 6 presents results of studies conducted at near-room temperatures (partially adopted from Lowson, 1982). In most studies the rate is not pH dependent, as was explained in Section 1.3.1. If one compare these rates to the oxidation rates in Table 5 one can see that the Fe(II) oxidation at pH<2 is very slow. However, some possible catalysts are reported in literature. The following solids and solutes were found to have a catalytic effect: in near neutral and neutral waters - alumina, silica, bentonite and ferric hydroxide precipitates; in acid solutions - the surface catalysts palladium, platinum, gold, and coconut charcoal and the solution catalyst Cu2+, Co2+, PO42-, P2O72- and organic chelating agents (Stumm and Lee, 1961; Lowson, 1982; Houben, 2004). The anions PO42- or P2O72- increase the oxidation rate of Fe(II) by O2 due to complexing of the ferric ion product (Cher and Davidsson, 1955; King and Davidsson, 1958). The anions F- and Cl- show similar effects (King and Davidson, 1958; Rönnholm et al., 1999). In general, the rate increases as the complexing affinity of the anion for ferric ion increases (Huffman and Davidson, 1956; King and Davidson, 1958; Stumm and Morgan, 1996). Thus, at a given pH, the rate is found to decrease in the series pyrophosphate, 41 phosphate, chloride, sulfate and perchlorate. It is to be expected that the complexing affinity for ferrous ion would be of the same order. The reported rate law is − d (Fe( II ) ) dt = 4k (Fe( II ) )PO2 for the first two media listed above and dt = 4k (Fe( II ) ) PO2 for the last three. It has been reported that the reaction is − d (Fe( II ) ) 2 also bimolecular with respect to Fe(II) and O2 in the presence of fluoride (Huffman and Davidson, 1956). Another possible catalyst is Cu2+. It is well known that addition of trace quantities of Cu2+ can catalyze the oxidation of dissolved Fe(II) by O2 in acid to neutral pH (Cher and Davidson, 1955; Huffman and Davidson, 1956; Stumm and Lee,1961; Lowson, 1982; Chmielewski and Charewicz, 1984; Zhang et al., 2000). In conclusion, the oxidation rate of ferrous iron by oxygen in natural waters is rapid due to the presence of hydroxyl ions and the carbonate system, which form complexes with Fe(II) that oxidize rapidly. At low pH levels the main species is the free ferrous ion that has a very low oxidation rate. However, the anions in solution may have a significant effect on the oxidation rate. In general, the rate increases as the complexing affinity of the anion for ferric ion increases. Pyrophosphate and phosphate anions have the highest acceleration effect. A catalytic effect by Cu2+ is also reported to be significant. 42 Table 6. Survey of reported Fe(II) oxidation rates under various conditions (Source: Lowson, 1982) Maximal Rate [M/s]* Rate equation 2 kt х [Fe(II)] х pO2 + 8.68E-11 2+ 2+ [Fe2+] Rate constant -5 -1 -1 -1 kt=1.5×10 M s atm -3 -1 -1 39 kCu х [Fe ][Cu ] kCu=7.6×10 M s 3.65E-10 k х [Fe2+]2 х pO2 2.78×10-6 M-1s-1atm-1 5.37E-6 1.23E-8 1.83E-8 3.36E-8 3.78E-8 4.68E-8 k1+k2 х [Cl-] k х [Fe2+]2 х pO2 k х [Fe2+]2 х pO2 k = [H+]-0.23 k х [Fe2+]2 х pO2 k х [Fe2+]2 х pO2 k х [Fe2+]2 х pO2 1.90E-7 0.001 M [media] Media 0.5 M 1.10E-05 M SO4 pH T [0c] Reference ~1 30 Huffman and Davidson, 1956 2- 2+ Cu 1M H2SO4 ~0 30 k1=2.7×10-6, k2=8.9×10-7 6.5×10-7 M-1s-1atm-1 0.001-0.025 M 0.2 M 0.01-0.3 M 0-3 M 1N HCl HCl <0 ~0 30 30 Huffman and Davidson, 1956 Matseevskii, 1980 McBain, 1901 9.7×10-7 M-1s-1atm-1 0.01-0.3 M 1N HClO4 ~0 30 George, 1954 4.0×10-6 M-1s-1atm-1 2×10-6 M-1s-1atm-1 1.1×10-6 M-1s-1atm-1 0.15-0.2 M 0.01-0.3 M 0.02-0.45 M 1M 1N 1M H2SO4 HClO4 H2SO4 ~0 ~0 ~0 30 30 30 kх[Fe2+]2х pO2 х[H+]-0.25 5.1×10-7 M-.0.75s-1atm-1 0.01-1.0 M 1N H2SO4 ~1 30 kх[Fe2+]х pO2 х[H3PO4]2 k1х[Fe2+]хpO2 х[H3PO4]2 + 7.88E-6 2+ k2х[Fe ]хpO2х[H4P2O7] kх[Fe2+]хpO2хf{HCl} kх[Fe2+]хpO2хf{HCl} * mg Fe(II)/l/h = 2.0х108 M/s 1.25×10-3 M-2s-1atm-1 0.005-0.02 M 1N H3PO4 H3PO4 + H4P2O7 HCl HCl ~2 30 Lamb and Elder, 1931 Chernaya, 1980 McBain, 1901 Mathews and Robins, 1972 Cher and Davidson, 1955 1-2 30 King and Davidson, 1958 <1 <1 18 45 Posner, 1953 Iwai et al, 1979 5.83E-7 k1=1.08×10-3 M-2s-1atm-1 k2=2.13×10-3 M-1s-1atm-1 1.3×10-3 s-1atm-1 3.1×10-4 s-1atm-1 0.005-0.02 M 0.055 N 0.2 M 43 0.2-0.4 M 0.2-0.8 M 8N 6N 6.4. Electrochemical oxidation of Fe(II) The use of electrolysis to oxidize ferrous iron has been investigated mainly in association with treatment processes for acid mine drainage (AMD) (Bunce et al., 2001) and disinfection of drinking water (Khelifa et al., 2004). The regeneration of ferric iron in an electrolytic cell can be achieved by direct or indirect oxidation. Direct oxidation is referred as the oxidation of Fe(II) to Fe(III) on the anode surface, whereas indirect oxidation is accomplished by producing a strong oxidizing agent through electrolysis that is subsequently used to oxidize the Fe(II). 6.4.1. Direct oxidation of Fe(II) The direct electro-oxidation process is based on oxidation of Fe(II) to Fe(III) on the anode surface (Eq. (31)) and evolution of hydrogen gas on the cathode surface (Eq. (32)) (Bunce et al., 2001). Fe 3+ + e − → Fe 2+ anode EH0=0.769 V (31) 2 H + + 2e − → H 2 cathode EH0=0.092 V (32) A competing (unwanted) reaction that invariably occurs on the anode is the oxidation of water to dissolved oxygen: 1 2 O 2 + 2 H + + 2e − → H 2 O anode EH0=1.226 V (33) Additionally, Ferric ions formed at the anode and those that are initially present in the solution can be back reduced at the cathode (Bisang, 2000; Bunce et al., 2001). 6.4.2. Indirect oxidation of Fe(II) The electro-generation of strong oxidizing agents, especially chlorine and/or hypochlorite for the removal of low concentrations of organic matter is a process that is known for many years (Alvarez-Gallegos and Pletcher, 1998). When elemental chlorine (Cl2(g)) is introduced into the water at natural pH it rapidly reacts with H2O (Eq. (34)) to form hypochlorous acid: Cl 2 ( g ) + H 2 O → HOCl + H + + Cl − pK=3.40 (34) Hypochlorous acid exists in solution in equilibrium with hypochlorite ions (OCl-) according to pH (pKa= 7.46), as shown in Eq. (35). HOCl ↔ OCl − + H + pK a = 7.46 44 pKa = 7.46 (35) These three oxidizing forms are denoted Active chlorine or AC in short. At the pH range 3-10.5 the hydrolysis rate (Eq. (34)) is rapid and almost constant. At pH values above 10.5 the rate sharply increases with an increase in pH (Spalding, 1962). However, at strongly acidic solutions (pH<3) chlorine hydrolysis is negligible (Spalding, 1962) and the dominant species in the aqueous phase is Cl2(aq). Since transporting, storing and handling of liquid chlorine are dangerous and ecologically unsafe (Kelsall, 1984; Krstajić et al., 1987) investigation and development of local, on-site processes for active chlorine production has gained increasing interest. The major process for an on-site AC production is electrolytic conversion of chloride ions in sodium chloride (NaCl) solutions into desired hypochlorite or chlorine gas as in Eq. (36). Cl 2 ( g ) + 2e − → 2Cl − anode EH0=1.392 V (36) The oxidation of ferrous iron by AC has been investigated mainly in the context of disinfection for drinking water (Khelifa et al., 2004). The products of Fe(II) oxidation are Fe(III) and chloride ions: or Cl 2 + 2 Fe 2+ → 2Cl − + 2 Fe 3+ (37) HOCl + 2 Fe 2+ + H + → Cl − + 2 Fe 3+ + H 2 O (38) 6.4.2.1. Competing reactions Several competing reactions for the electro-generation of active chlorine have been reported, contributing to cell efficiency reduction and byproduct formation, (Krstajić et al., 1991; Rudolf et al., 1995): At the anode hypochlorite may be oxidized to chlorate according to Eq. (39). − 6OCl − + 3H 2 O → 2ClO3 + 4Cl − + 6 H + + 12 O2 + 6e − anode EH0=1.446 V (39) Another competing reaction at the anode is oxygen formation (Eq. (33)). Hypochlorite may be also reduced at the cathode to the chloride ion form as described in Eq. (40) or chemically transformed into chlorate or chloride according to Eq. (41) and Eq. (42) respectively: OCl − + H 2 O + 2e − → Cl − + 2OH − cathode EH0=1.481 V (40) − 2 HOCl + OCl − → ClO3 + 2Cl − + 2 H + (41) 2OCl − → 2Cl − + O2 (42) 45 The ratio between the anode surface area to the cathode surface area (Sa:Sc) may affect the rate of chlorate production on the anode and its reduction on the cathode. Khelifa et al., (2004) reported that the optimum Sa/Sc value in their experimental system was 1.33 for anode made of Ti/TiO2 and cathode made of Ti. Nevertheless, the rate of chlorate formation according to Eq. (41) depends on the temperature and is higher at elevated temperatures. At relatively low temperatures (about 20-25 °C) chlorate formation is slow (Grinberg et al., 2001). The pH of solution also affects chlorate formation. First of all Cl2 hydrolysis to HOCl is relatively slow below pH3 (see above), and since hypochlorite is the reactant for chlorate production - chlorate formation should be minor below pH3. Secondly, the main reaction on the cathode around pH2 was reported to be H2 evolution (AlvarezGallegos and Pletcher, 1998; Bisang, 2000) (Eq. (32)). 6.4.2.2. Oxidation of Fe(II) by Cl2 The following mechanism was proposed by Gilliland et al. (1958) and confirmed by Hikita et al. (1975): step I: Cl 2 + Fe 2+ ↔ Cl 2 ⋅ Fe 2+ (43) step II: Cl 2 ⋅ Fe 2+ + Fe 2+ ↔ 2Cl − + 2 Fe 3+ (44) The first step is an irreversible second-order reaction (first order with respect to both ferrous iron concentration and chlorine concentration) and is the rate controlling step of the overall reaction. The second step is expected to be second order (Gilliland et al., 1958). The information on oxidation of Fe(II) by Cl2 is scarce. Yet, the electro-oxidation of chloride ions to active chlorine is prevalent in literature. 6.4.3. Electrolytic production processes of active chlorine The production of active chlorine depends on various operational parameters; some of these parameters are discussed hereunder. 6.4.3.1. Divided or undivided cell Electrolytic cells appear in the form of divided or undivided cells. In an undivided cell, the cathode and anode electrodes are merged in the same solution, without any 46 separation of the electrolyte solution. A divided cell, on the contrary, is divided into compartments – one for each electrode, which make two different electrolyte solutions: an anolyte solution for the anode and a catholyte solution for the cathode (Bunce et al., 2001). Divided cells produce more concentrated hypochlorite solutions and have higher current efficiency for hypochlorite production as compared with undivided cell processes (current efficiency is the ratio of the electrochemical equivalent current density for a specific reaction to the total applied current density). The downside of this method is that electrolysis in membrane-divided-cells for hypochlorite and chlorine production is susceptible to failure due to clogging and formation of precipitates on the membrane surface (Bisang, 2000). This affects membrane permeability and, consequently, overall process efficiency. An undivided cell is prone to current efficiency loss due to the unwanted reactions and byproducts formation, out of which chlorate formation is probably the most problematic. However, as mentioned earlier, working at room temperature and at low pH decreases remarkably the production of chlorate. 6.4.3.2. Batch or flow-through mode Electrolytic cells can be operated in batch mode where the solution is continuously mixed during electrolysis (Krstajić et al., 1987; Khelifa et al., 2004). Another approach is flow-through electrolysis where two vessels are used, one as the solution carrier and another as the electrolytic cell. The electrolyte is recycled from the holding vessel to the electrolytic cell (Rudolf et al., 1995; Kraft et al., 1999; Grinberg et al., 2001). The flow-through cell can be operated as a single pass mode (Kelsall, 1984; Grinberg et al., 2001) or in a recirculation mode (Robertson et al., 1983; Rudlof et al., 1995). Flow-through systems are favorable because the forced convection prevents local overheating of the electrolyte near the electrode surface (Grinberg et al., 2001), reduces gas locking of the cathode, provides similar current density distribution across the electrodes area and prevents the precipitation on the electrodes (Kelsall, 1984). Other operational parameters to be considered are: flow-through velocity, chloride concentration, anode and cathode materials, interelectrode gap (i.e., the distance between the anode and the cathode) and current density. 47 In Table 7 the conditions and results of three experiments for hypochlorite production, all in undivided cells, are summarized. The conditions of the experiments presented in Table 7 vary, but one can see that chlorate formation rates of 8 g/h can be achieved, which means that similar rates are possible for Fe(II) oxidation. Table 7. Conditions and results of experiments on chloride electrochemical oxidation in an undivided cell Rudolf et al., 1995 Cell type Flow-through rate Initial NaCl concentration Electrolyte temperature and pH Grinberg et al., 2001 Khelifa et al., 2004 Flow-through, Flow-through, single recirculating pass 5.16 l/min 0.20 l/min - 15 and 30 g/l 8.9 g/l 3M 25°C initial pH=6 20°C batch 20°C pH>8 Titanium covered Anode material with 30 mol% Ti/Pt Ti/RuO2 Stainless steel Ti/Pt Titanium 1 0.92 1.33 4 mm 0.5 mm 5 mm 0.12 kA/m2 3.5 kA/m2 RuO2 and 70 mol% TiO2 Cathode material Sa/Sc ratio Interelectrode gap Current 1.059 – 4.237 densities kA/m2 sodium Calculated hypochlorite: 7.8- oxidation rate 10.0 g/h / chlorate: 0.3-0.9 g/h active chlorine: 8.2 sodium hypochlorite: g/h / the only active 11.0 g/h / chlorate: chlorine species <0.2 g/h detected was hypochlorite 48 In summary, it is possible to use an electrolytic process to oxidize Fe(II), directly or indirectly (i.e., in the presence of chlorides). The avoidance of competing reactions should be considered. Many operational parameters can be optimized in order to achieve a high oxidation rate of Fe(II). In the current work the oxidation of Fe(II) by O2 in the presence of catalysts and electrochemical oxidation of Fe(II) were investigated as part the development of an LRSR process at low pH. 49 7. Hypothesis and objectives 7.1. Research hypothesis It is possible to develop an LRSR process that is operated at low pH (pH<1.5), under such conditions where: (1) precipitation is insignificant and does not impede the process, (2) reactive-absorption of H2S(g) is efficient and (3) Fe(III) regeneration (i.e. Fe(II) oxidation) rate is equal or higher than the rate of H2S(aq) oxidation by Fe(III). Two process options were investigated with respect to the research hypothesis: LRSR process at pH1.0 with catalytic oxidation of Fe(II) by atmospheric O2 This option is based on the use of Cu2+ and H3PO4 which have been reported in the literature to accelerate the Fe(II) oxidation rate at low pH. A bubble column reactor is filled with solution that contains high concentration of Fe(III) (9 g Fe/l) and copper and/or phosphate. A gaseous stream contaminated with H2S(g) is introduced into the reactor. The concentration of Fe(III) is maintained high in order to obtain a high reactive absorption efficiency. The Fe(II), generated as a result of H2S oxidation by Fe(III), is oxidized back to Fe(III) by O2, which is introduced into the reactor with the gaseous stream. If the O2 content in the gaseous stream is low, air or oxygen should be supplied to the reactor. The Fe(II) oxidation rate in the presence of copper and/or phosphate should be high enough in order to maintain the Fe(III) concentration high and steady. LRSR process at pH1.0 with electrochemical oxidation of Fe(II) This option is based on the use of an electrolytic cell to accelerate the Fe(II) oxidation rate. A bubble column reactor is filled with solution that contains a high concentration of Fe(III) (9 g Fe/l). The oxidation of the Fe(II) can be preformed directly or indirectly. In the case of indirect oxidation chloride ions are also added to the solution. A gaseous stream contaminated with H2S(g) is introduced into the reactor. When the concentration of Fe(II) generated exceeds a certain level (2 g Fe/l), a portion of the solution is recycled into a flow-through electrolytic cell. In the direct oxidation option the Fe(II) is oxidized on the anode surface, while in indirect oxidation option Cl- is oxidized on the anode surface to Cl2 that oxidizes the Fe(II). When the concentration of Fe(II) goes down below a certain level (1 g Fe/l) the circulation into the electrolytic cell is stopped. Such an operation enables maintaining a relatively constant and high Fe(III) concentration, while the fact that Fe(II) 50 oxidation is not operated to very low Fe(II) concentrations has a kinetically advantage. The use of a bubble column reactor has the advantage of a high contact area between the gas and the solution which suppose to reduce the size of the scrubbing reactor. Apart from that, the type of the reactor that is used is of no importance to the process. 7.2. Research objectives 1) Assessment of the potential of long-term formation of Fe(III) precipitates from the H2S reactive-absorption solution as a function of pH and the composition of solution. 2) Investigation of the effect of pH and H3PO4, Cu(II) and Cl- concentrations on the reactive-absorption efficiency of H2S(g). 3) Investigation of the effect of Fe(III), H3PO4 and Cu(II) concentrations on Fe(II) oxidation by O2 at pH1.0. 4) Investigation of the effect of Cl- concentration and current density on Fe(II) electrochemical oxidation. 5) Determining the most advantageous conditions (or range of conditions) for catalytic oxidation of Fe(II) at pH1.0, with respect to H2S(g) removal by the LRSR process, and evaluation of the Fe(II) oxidation rate. 6) Determining the most advantageous conditions (or range of conditions) for electrochemical oxidation of Fe(II) at pH1.0, with respect to H2S(g) removal by the LRSR process, and evaluation of the Fe(II) oxidation rate. 51 8. Materials and Methods 8.1. Chemicals All reagents were analytical grade if not stated otherwise. Table 8. Description of the reagents used in experiments Name of chemical Formula Acetic acid glacial CH3COOH Ammonium acetate CH3COONH4 Ammonium iron(II) sulfate 6-hydrate (FAS) Copper sulfate Manufacture Remarks BIO-LAB ltd. 99.8% LOBA CHEMIE (NH4)2Fe(SO4)2•6H2O Merck CuSO4 Merck Hydrochloric acid HCl Frutarom Hydrogen sulphide H2 S BOC Iron(II) sulphate 7hydrate Iron(III) sulfate exsiccated FeSO4•7H2O Fe2(SO4)3 Sigma-Aldrich C.P. 85% H3PO4 Carlo Erba Sodium chloride NaCl Frutarom Sodium fluoride NaF Riedel de Haen Sodium hydroxide NaOH Frutarom Sulfuric acid H2SO4 Frutarom monohydrate 5-Sulphosalicylic acid 2hydrate Nitric acid Mercury thiocyanate C12H8N2•H2O C7H6O6S•2H2O Min. 98% Fluka Riedel de Haen HNO3 Merck C2HgN2S2 Fluka 52 Background Carlo Erba ortho-Phosphoric acid 1,10-Phenanthroline 500 ppm, N2 65% 8.2. Analytical equipment 1) Filter paper, 0.45 µm, White 47 mm Gridded, Sterile, Millipore 2) (ICP) emission spectrometer, Optima 3000 DV, Perkin Elmer 3) Kitagawa gas detector tube system for H2S(g) (1-60 ppm, 0.5-40 ppm and 0.2-6.0 pH meter, Metrohm 827 pH lab, with Metrohm 6.0228.010 glass electrode 4) Linear shaker with the water bath, BT-350, MRC Laboratory equipment 5) MINEQL+ software version 4.5 (Envrionmental Research Software, Hallowell, ME) 6) QuickChem 8500 (LACHAT instruments, USA) for ammonia detection 7) Spectrophotometer, Genesys 10, Spectronics 8) Syringe Driven Filter Unit, 0.22 µm. Millex-GV, Millipore 9) ppm), Komyo Rikagaku Kogyo K.K. 10) Tubes gas dispersion pyrex fritted disc 40-60 microns 11) 0-20A DC power supply device (HY3020, Aviv Energy tech. Ltd, Israel) 8.3. Experimental All experiments were conducted at room temperature (25 ±1 °C); pH adjustments were done by the addition of concentrated (98%) sulfuric acid or NaOH 5N. The applicability of each one of the proposed processes was investigated according to the following parameters: (1) H2S(g) reactive-absorption efficiency, (2) practical precipitation potential of the working solution and (3) Fe(II) oxidation rate. 8.3.1. H2S(g) reactive-absorption experiments The efficiency of H2S(g) reactive-absorption into acidic solutions containing Fe(III) was studied using a bubble column reactor. Mixtures of hydrogen sulfide and air were bubbled through a bubble column reactor filled with acidic ferric sulfate solution. The flow rates of both the air and the hydrogen sulfide gas were measured and controlled. The total flow rate was kept low enough in order to maintain a homogeneous bubble flow. The desired concentrations of H2S and mixed air flow rates were achieved by controlling the individual flow rates of the N2-H2S mixture and air. The concentration of H2S in the experiments did not exceed 80 ppm, mainly for reasons of safety. The 53 H2S(g) concentrations at the inflow and the outflow of the reactor were measured by Kitagawa gas detector tubes designed for H2S(g). Each measurement was repeated two or three times. 8.3.2. Determination of practical precipitation potential of the working solutions Solutions with different compositions were left for several weeks to months in order to follow the formation of precipitates. Samples from solutions were taken and filtered every few weeks and the concentration of several dissolved components (e.g., iron, copper) was measured by ICP analysis or via a spectrophotometric method. Samples of the solid phase were also taken few times. The samples were dried in an oven and then dissolved in order to determine the presence of part of the components and their concentrations ratio. 8.3.3. Determination of Fe(II) oxidation rate 8.3.3.1. Catalytic oxidation of Fe(II) Batch experiments were conducted with initial ferrous iron concentration of 1000 mg Fe/l (from (NH4)2Fe(SO4)2•6H2O) in 500 ml bottles. Air was bubbled to supply oxygen through gas dispersion tubes. Fig. 12 is a photo of the system that was used for the catalytic oxidation experiments. The dissolved oxygen concentration was measured and verified to be the saturation concentration. Every few hours the bottles were weighed and refilled to the initial weight with diluted H2SO4 solution (~0.1 N). Samples from solutions were taken and diluted in diluted H2SO4 solution in order to determine the dissolved Fe(II) concentration. The concentration of Fe(II) was measured by the modified Phenanthroline method developed by Herrera et al. (1989). Each sample was measured three times, and if the standard deviation of the three samples was less than 5% the average value of the three was taken. Otherwise only two replicates were considered. 54 Figure 12. A photo of the experimental system for catalytic Fe(II) oxidation 8.3.3.2. Electrochemical oxidation of Fe(II) The electrochemical laboratory system setup that was used is shown in Fig. 13. A flow-through electrolytic cell with parallel rectangular electrodes was used. The electrolyte was recycled between the flow-through cell and the electrolyte holding vessel by a peristaltic pump. A PVC pipe reactor, (internal diameter = 5.71 cm) was used as a holding vessel to provide the Cl2(g) formed on the anode a longer retention time inside the system. A 0-20A DC power supply device was used in order to provide the electrical current. Three cathode dimensions were used: 0.72×9.15 cm2, 3.0×9.14 cm2 and 5.1×9.15 cm2. Anode dimensions were 5.1×9.15 cm2 in all experiments. Electrode materials, inter-electrode gap and flow-through velocity were adopted from literature related to the electrochemical hypochlorite production, as discussed hereunder. The most effective anode material for hypochlorite production was reported to be titanium coated by ruthenium oxide (Ti/RuO2) due to its low over-potential for chloride oxidation and mechanical and chemical stability (Ponzano, 2007). Cathode material may affect the hypochlorite back reduction rate and hydrogen evolution rate (Khlaifa et al., 2004). Ponzano (2007) recommended on titanium coated with iridium 55 oxide and plain titanium as good cathode materials, while graphite, stainless steel and nickel were found less suitable. In the current study titanium electrodes were used. Khelifa et al., (2004) studying a batch electrolyzer for hypochlorite production and reported the optimal interelectrode gap to be <8 mm. In other studies the interelectrode gap used was 4 mm or shorter (Kelsall, 1984; Robertson et al., 1983; Rudolf et al., 1995). An inter-electrode gap of 3.46 mm was used in the current work. The oxidation of hypochlorite to chlorate on the anode depends on the hypochlorite concentration in the vicinity of the anode (Grinberg et al., 2001). Thus the flowthrough velocity may affect the rate of chlorate production. The flow-through velocities that were found in literature were 26.275 cm/s (Rudolf et al., 1995) and 0.27-1.08 cm/s (Grinberg et al., 2001). The electrolyte recirculation rate between the electrolysis cell and the holding vessel in the current work was 1.72 l/min and consequently the flow velocity in the electrolysis cell was 16.24 cm/s. Figure 13. A scheme (on the left) and a photo of the laboratory system for Fe(II) electrochemical oxidation experiments. 1- Electrolyte holding (Cl2 absorption) vessel; 2- Flow-through electrolysis cell; 3- DC power supply; 4- Peristaltic pump; 5Chlorine trap. In the direct oxidation experiments (Cl-=0) the concentration of dissolved Fe(II) was detected by measuring the oxidation-reduction potential (ORP) at predetermined time 56 intervals. In the indirect oxidation experiments the concentration of dissolved Fe(II) was determined using the modified Phenanthroline method (Herrera et al., 1989). All the gas that was formed during the electrolysis was passed through a measuring glass (250 ml) filled with 1M NaOH solution. At this alkaline solution Cl2(g) that formed on the anode surface and did not react with Fe(II) was hydrolyzed according to Eq. (34). Consequently, the determination of the chloride ions concentration in the trap solution was used to assess the chlorine mass that potentially escaped the system during the operation. 8.4. Analytical methods 8.4.1. Determination of the total dissolved iron, phosphate and copper concentrations The concentrations of the total dissolved iron, phosphate and copper were analyzed by inductively coupled plasma (ICP) emission spectrometry, Optima 3000 DV, Perkin Elmer. 8.4.2. Determination of dissolved ferrous iron concentration The concentration of dissolved ferrous iron was measured spectrophotometrically according to the modified phenanthroline method (Herrera et al., 1989). The method is highly accurate (below 5% error) in the presence of ferric iron if the Fe(II) percentage is higher than 5% of the total dissolved iron. Samples were diluted with diluted H2SO4 solution (pH1.0) in order to lower the Fe(III) concentration below 200 mg Fe/l. Two modifications were made to the method: the absorbance of a sample was read 10 minutes after the addition of the colorimetric reagent instead of 5 minutes and diluted H2SO4 solution (pH1.0) was used instead of distilled water. 57 8.4.3. Determination of dissolved ferric iron concentration The concentration of dissolved ferric iron was measured via a spectrophotometric method using Sulphosalicylic acid (Zolotov, 2001). Samples were diluted with diluted H2SO4 solution (pH1.0) in order to lower the Fe(III) concentration below 20 mg Fe/l. 8.4.4. Determination of chloride concentration The concentration of chlorides was measured via a spectrophotometric method using mercury thiocyanate and iron(III) alum (Yoshinaga and Ohta, 1990). Yoshinaga and Ohta (1990) reported that higher accuracy was obtained with higher concentrations of mercury thiocyanate and FAS reagents. Thus, the concentrations used in current work were the highest reported by Yoshinaga and Ohta (1990), i.e., 9 g/l mercury thiocyanate and 300 g/l of FAS. The method is reliable also in highly basic solutions (1M NaOH), although a specific calibration curve should be made. 8.4.5. Determination of ammonia concentration The concentration of ammonia was determined by flow injection analysis of QuickChem (method 10-107-06-2-A). The ammonia concentration is determined by a spectrophotometric measurement, based on the color that is produced when ammonia is heated with salicylate and hypochlorite in an alkaline phosphate buffer. 8.4.6. Analysis of species distribution by the MINEQL+ software Input data - components of solution, total concentrations of the components and pH. Ionic strength was calculated during the runs of the software. Output data – concentrations of each one of the species of each component. Data on precipitation potential of all possible solids is also given. It should be mentioned that MINEQL+ analysis is based on thermodynamic data and kinetics are not considered. In addition, ionic strength corrections in MINEQL+ are calculated according to the Debye-Huckel equation, which means that the accuracy of thermodynamic constants is lower at ionic strengths exceeding 0.5 M. In the current 58 study the ionic strength of most of the solutions exceeded 0.5M, thus results should be considered to be semi-qualitative rather than quantitative. 59 9. Results and Discussion The Results chapter is divided in two according to the two processes that were investigated in the work: LRSR process at pH1.0 with catalytic oxidation by O2 of Fe(II) and LRSR process at pH1.0 with electrochemical oxidation of Fe(II). In each section the results of the following three subjects are presented: (1) reactiveabsorption efficiency of H2S(g), (2) practical precipitation potential of Fe(III) species in the working solution and (3) Fe(II) oxidation rate, all under various operational conditions. 9.1. Catalytic oxidation The compositions of the solutions that were tested in the catalytic oxidation experiments were: Fe(III) concentrations: 0, 81, 161 and 322 mM = 0, 4.5, 9.0 and 18.0 g Fe(III)/l Total-phosphate concentrations: 0 to 867 mM = 0 to 26.7 g P/l Cu(II) concentration: 0 to 8 mM = 0 to 0.5 g Cu(II)/l. The initial Fe(II) concentration was 18 mM (= 1 g Fe(II)/l) in all experiments. 4.2.4. H2S(g) reactive-absorption efficiency The reactive-absorption efficiency was tested once in a solution with the following composition: 18 mM Fe(II), 81 mM Fe(III), 220 mM H3PO4 and 2 mM Cu(II) at ( ) 3 3 and it contained around 80 ppm pH1.22. The flux of gas was 1.25 mair / min / m solution of H2S(g). The reactor had an internal diameter of 5.71 cm, and air was sparged by a gas dispersion tube. In less than one hour a colloid-type dark precipitate appeared in the solution. The precipitate was identified as the mineral Covellite (CuS). Covellite is described as thin crusts or as deep blue-black powdery or sooty masses (Dunn and Muzenda, 2001). The structure of CuS covellite is best represented as Cu(I)4(S-I)4Cu(II)2S(-II)2 (Madarász et al., 2001). The precipitation of CuS was corroborated also by the measurements of dissolved copper concentration. As shown in Fig. 14 the concentration of dissolved copper decreased linearly with time, which confirmed the CuS precipitation observation. The reactive-absorption efficiency of H2S(g) with in the experiment time is also presented in Fig. 14. The reactive-absorption efficiency was high at the beginning of 60 the experiment (80-85%), and then decreased. This is probably due to the decrease in the availability of Cu(II), which precipitated with S2- to form CuS. This means that the oxidation of H2S(aq) in the tested solution was relatively slow and that precipitation of CuS was the main mechanism by which H2S(g) was removed. Zhang and Tong (2006) claimed that once CuS forms, it is rapidly oxidized by Fe(III) to S0 (see1.2.4.2). This is in contrast with the results presented here. It is possible that the higher concentrations of Fe(III), Fe(II) and Cu(II) and the different concentrations ratio, that were applied by Zhang and Tong (2006) (see Subsection 1.2.4.2), are the reason for this difference in results. However, the very high concentrations that were suggested by Zhang and Tong (2006) are problematic for blow-down reasons. The solution contains high concentration of copper, whose release to the environment is limited. In addition, Zhang and Tong (2006) reported that no precipitation (other than elemental sulfur) was observed after 4 hours of experiment. Yet, it is very unlikely that in such a concentrated solution precipitation will not occur in long-term [%] H2S(g) reactive-absorption efficiency operations (which are much longer than the 4 hours they reported on). 0.90 1.6 0.85 1.4 0.80 1.2 0.75 1.0 0.70 0.8 0.65 0.6 y = -0.2352x + 1.4287 R2 = 0.993 0.60 0.4 0.55 0.2 0.50 0.0 0 1 2 3 Time [h] 4 5 6 Figure 14. Change in H2S(g) reactive-absorption efficiency (◊) and dissolved copper concentration (х) with time. 61 In conclusion, the addition of Cu(II) into a solution that is used for an LRSR process is problematic because at the conditions tested, because Cu2+ apparently precipitates with S2- to form CuS. The addition of H3PO4 as a catalyst was also found problematic, as discussed in the following section. 4.2.5. Practical precipitation potential in the working solution Precipitation was observed in all Fe(II) oxidation experiments (Section 4.1.3) in which phosphate concentration in solution was 0.16 M (4.6 g P/l) or higher. It took between a few hours to a few days for the precipitates to become visible. Two types of precipitates were observed: the first was bright-yellowish and powder like; and the second bright-purplish and crust like (Fig. 15). Samples from 5 different bottles were filtered (after standing untouched for 170 days) and the solid phase was dried in an oven at 60°C for 24 hours. The dried phase was then dissolved in highly acid solution (HCl 1N) and the samples were taken for ICP analysis, in order to quantify the content of iron, phosphate and copper in the solid. In all 5 samples the concentration of copper was below the detection limit. The concentrations of total-phosphate (PT) and total-iron (FeT) found in the solid phase are presented in Table 9. In all 5 samples the molar ratio of phosphate to iron was found to be between 1 and 2. The ICP analysis of the solutions from other Fe(II) oxidation experiments showed a clear decrease in the concentration of phosphate and iron, relative to the initial concentrations, in most of the samples. A decrease in ammonia concentration was also observed, probably due to physical adsorption of ammonia on the solid surface. Figure 15. A photo of precipitates from catalytic oxidation experiments after 170 days 62 Table 9. Composition of precipitates in solutions from the catalytic oxidation experiments (after 170 days) Total phosphate and total iron concentrations Sample FeT FeT Color PT [mg/l] 2 yellow 9.8 ± 0.6 16.5 ± 0.1 0.32 0.30 1.1 3 yellow 55.7 ± 0.7 75.0 ± 0.2 1.80 1.34 1.3 4 purple 16.9 ± 0.1 26.2 ± 0.8 0.54 0.47 1.2 5 purple 16.8 ± 0.2 25.6 ± 0.6 0.54 0.46 1.2 6 purple 26.5 ± 1.6 25.8 ± 0.0 0.86 0.46 1.9 number [mg/l] PT [mM] PT/FeT ratio [mM] M/M The following solution was analyzed by MINEQL+ in order to evaluate its thermodynamic precipitation potential and compare it to the observed results: 18 mM Fe(II), 81 mM Fe(III), 59 mM H3PO4 and 0.8 mM Cu(II) at pH1.0. The precipitation potential was found positive for the minerals Strengite (FePO4•2H2O) and H-Jarosite (HFe3(OH)6(SO4)2). In fact, 99% of total-H3PO4 was expected to precipitate as strengite at steady state, and the rest of the Fe(III) was projected to precipitate as jarosite. This is equivalent to 20.6 g of solids for 1 l solution, which is much higher than the amounts of solid that were visually observed to precipitate. In the actual experiments, which lasted for 3 months it was assumed that at pH conditions as low as pH1.0 the kinetics of precipitation would be very slow and thus solid precipitation would appear after a very long time (in the order of months). If this was indeed the case, the loss of Fe(III) due to precipitation would be tolerable, from an engineering perspective. The assumption of slow kinetics at this pH was supported by Cher and Davidson (1955), who did not report on any precipitation in their experiments, which were conducted at pH 1.1-1.7 and with PT between 0.402 and 1.0 M. However, in their experiments the initial Fe(II) concentration was 10 mM and no Fe(III) was added, which may explain the fact that precipitation was not observed: even if all the Fe(II) was oxidized to Fe(III), the maximal precipitation possible is of 0.56 g Fe per 1 l of solution.. 63 All in all, the addition of the potential catalysts Cu2+ and H3PO4 to an LRSR solution was found to be inapplicable in the range of conditions that were tested in the work. Nevertheless, the investigation of the effect of Cu2+ and H3PO4 on Fe(II) oxidation by O2 at low pH was continued because of its scientific value. 4.2.6. Fe(II) oxidation rate in the catalytic oxidation experiments First, the effect of phosphate or copper on the modified phananthroline method (Herrera et al., 1989) was assessed. Calibration curves were prepared with different concentrations of H3PO4 and Cu(II). The different calibration curves that were attained were similar to each other and no distinguishable effect of H3PO4 and Cu(II) was observed. Consequently, a general calibration equation was chosen to represent all the solutions, as follows: y = 122.4 x − 1.902 (45) where y is Fe(II) concentration in mg Fe/l and x is the absorbance in %. All experiments were conducted for 4 to 20 hours. In each experiment the concentration of Fe(II) was measured 6 to 12 times. Examples of the results of 4 different experiments are presented in Fig. 16. Fe(II) concentration [mg Fe/l] 1000 900 Fe(III)=81mM, Cu(II)=0, P=588mM Fe(III)=161mM, Cu(II)=0.8mM, P=588mM 800 Fe(III)=81mM, Cu(II)=16mM, P=588mM 700 Fe(III)=81mM, Cu(II)=0.8mM, P=588mM 600 500 0 1 2 3 Time [h] 4 5 6 Figure 16. Example for the results obtained in the catalytic oxidation experiments. Concentration of Fe(II) as a function of time in four different experiments. pH1.0 64 In most of the studies that addressed the spontaneous oxidation rate of Fe(II) by oxygen at low pH levels (pH < 2) the kinetic equation was found to be first-order with respect to oxygen, first or second-order with respect to Fe(II) and independent of pH. For the case that the rate is first order in Fe(II) the most widely accepted mechanism is the one presented by King and Davidson (1958): a ,b Fe( II ) + O2 + H + ←→ Fe( III ) + HO2 • • (46a,b) Fe( II ) + HO2 + H + → Fe( III ) + H 2 O2 (47) 2 Fe( II ) + H 2 O2 + 2 H + → Fe( III ) + 2 H 2 O (48) Equation (46a) is the rate determining step and Eq. (48) proceeds very rapidly. The reaction order with respect to Fe(II) concentration in the catalytic oxidation experiments was assessed under the assumption of zero-, first- and second-order kinetics. The best fit was used to calculate the oxidation rate according to 40% reaction completion, i.e. the time needed to oxidize 400 mg/l of Fe(II) (the initial concentration in all the experiments was around 1000 mg/l). Reducing the Fe(II) concentration below 400 mg/l has no advantage from the standpoint of the LRSR process. Second-order kinetics was found to be the most suitable for 40 out of 60 experiments. In 15 out of 60 experiments a zero-order dependency was the most fitting. No clear correlation could be observed between the reaction order (with respect to the Fe(II) concentration) and other components in the aqueous phase (i.e., iron, phosphate and copper concentrations). The effects of the following factors on the oxidation rate were investigated in the catalytic Fe(II) oxidation experiments: pH, initial Fe(III) concentration, total dissolved phosphate concentration (PT) and initial Cu(II) concentration. Since the counter anion of all the reagents was SO42- (and also H2SO4 was used to adjust the pH of the tested solutions) the effect of total sulfate concentration (ST) was also examined. 4.2.6.1. Effect of pH The pH was measured at the beginning and end of each experiment. The average change in pH was ±0.03 and maximal change was ±0.10. The pH was expected to increase with time, since Fe(II) oxidation consumes alkalinity (Eq. (19)). However, aqueous solution at such low pH levels has a high buffering capacity. Table 10 lists the conditions of experiments that were conducted at pH different then 1.0: the first four experiments were conducted at pH~1.6 and the rest – at pH~0.8. For each of 65 these experiments, a parallel experiment at pH1.0 was conducted (i.e. with the same initial concentrations of Fe(III), H3PO4 and Cu(II)). Table 10. The conditions of catalytic oxidation experiments that were conducted at pH values other than 1.0 and representative Fe(II) oxidation rates Concentration [mM] pH Representative Fe(II) oxidation rate [mg-Fe/l/h] PT ST Cu(II) Fe(III) A 613 31 7.9 0 1.60 684 B 807 31 7.9 0 1.55 888 C 807 62 0.0 0 1.59 57 D 613 94 0.0 0 1.64 28 E 743 281 0.8 161 0.78 25 F 291 312 0.8 161 0.83 9 G 452 312 0.8 161 0.80 14 H 581 343 0.8 161 0.75 18 I 0 499 0.8 161 0.77 6 J 387 530 7.9 322 0.78 4 K 387 530 0.0 322 0.77 5 L 613 530 7.9 322 0.72 10 M 807 530 7.9 322 0.67 16 A comparison of the representative ferrous oxidation rates (in the specified Fe(II) concentration interval) at different pH values is presented in Fig. 17 and Fig. 18. It is clear from Fig. 17 that at pH~1.6 the rate of Fe(II) oxidation is one order of magnitude higher than at pH 1.0. Although a higher oxidation rate is ostensibly advantageous, precipitation of Fe(III) oxides at pH~1.6 is rapid, making this pH infeasible for long term operation of the LRSR process. An attempt to conduct experiments similar to experiments A to D in the presence of Fe(III) (161 mM) failed, because a massive Fe(III) precipitation occurred (the solution was clear only when pH was lowered below pH1.3). 66 1000 684 888 pH>1.0 -d[Fe(II)]/dt [mg Fe/l/h] 242 pH=1.0 78 100 57 28 7 10 6 1 A B C Experiments D Figure 17. Comparison between rates of Fe(II) oxidation at pH~1.6 and pH~1.0 under various conditions (see Table 10 for specific operational conditions). y-axis is in logarithmic scale Since precipitation could not be avoided also at pH1.0 (see 4.1.2) in the presence of phosphate the oxidation rate was investigated also at lower pH (pH0.8, experiments E to M in Table 10). The effect of lowering the pH below 1.0 on the oxidation rate was not conclusive, as shown in Fig. 18. -d[Fe(II)]/dt [mg Fe/l/h] 35 pH=1.0 30 pH<1.0 25 20 15 10 5 0 E F G H I J Experiments K L M Figure 18. Comparison between rates of Fe(II) oxidation at pH~0.8 and pH~1.0 under various conditions (see Table 10 for specific operational conditions) Although the oxidation rates at pH0.8 were not much lower than the rates at pH1.0 it was decided to set the working pH at pH1.0. This was done for two reasons: first, the 67 precipitation was minimized but not avoided even at pH0.8, and second, the lower pH necessitates the use of even more resistant equipment, which makes the process less cost effective. All experiments presented from here onward were conducted at pH1.0 (±0.1). As mentioned earlier there was no apparent change in pH during the experiments. 4.2.6.2. Effect of initial concentration of Fe(III) An investigation of the H2S(g) reactive absorption efficiency in a bubble column reactor at low pH was reported only in one study, by Gendel (2007), who applied the following conditions: pH 1.72 and 81 or 161 mM Fe(III). Gendel (2007) found, like others before him (see Subsection 1.2.4.2), that an increase in the concentration of Fe(III) results in an increase in the reactive-absorption efficiency. Efficiencies higher than 95% and 75% were attained using Fe(III) concentrations of 161 and 81 mM respectively and air flux of 1 m3/m3/min. Since a bubble column reactor was used in the current study, the Fe(III) concentrations used by Gendel (2007) were adopted for the examination of their effect on the Fe(II) oxidation rate. Experiments with similar initial concentrations of Cu(II) and PT, were gathered into a series. In Fig. 19 the oxidation rates of 6 of these series are presented. The initial [Fe(III)]:[Fe(II)] ratio were 0, 4.5, 9 and 18. Error bars were added in order to emphasize the change in Fe(III) concentration during the experiments as a result of Fe(II) oxidation. It is clear from Fig. 19 that Fe(III) has an inhibiting effect on the Fe(II) oxidation rate (exponential tendency seems to explain well the retarding effect of Fe(III) concentration, 0.636<R2<0.998). From a chemical standpoint the retarding effect of Fe(III) may be explained by Eq. (46b), i.e. reduction of Fe(III) by HO2●. Pham and Waite (2008), who investigated the Fe(II) oxidation by O2 in natural waters, concluded that the reduction of Fe(III) is important particularly under conditions where precipitation of Fe(III) is minimal. This seems to explain well the decrease in Fe(II) oxidation rate that was obtained with an increase in the initial [Fe(III)]/[Fe(II)] ratio, because the experiments were conducted at pH1.0, where precipitation of Fe(III) is relatively low. 68 -d[Fe(II)]/dt [mg Fe/l/h] 250 [Cu]=0.8 mM [Cu]=1.6 mM [Cu]=3.1 mM [Cu]=6.3 mM [Cu]=7.9 mM [Cu]=15.7 mM 200 150 PT = 581 mM 100 50 0 0 50 100 150 200 250 300 350 Fe(III) initial concentration [mM] Figure 19. Fe(II) oxidation rate as a function of the initial Fe(III) concentration with various concentrations of Cu(II). Error bars present the change in Fe(III) concentration during the experiments Although the oxidation rates were relatively high when no Fe(III) was initially present, these results are meaningless with regard to the H2S(g) removal process since a high concentration of Fe(III) is essential for efficient reactive-absorption of H2S(g). From the standpoint of the LRSR process it might be more practical to use an Fe(III) concentration of 81 mM (4.5 g Fe/l) in the working solution rather than 161 mM (9 g/l), since the Fe(II) oxidation rate in the presence of 81 mM Fe(III) is almost twice as high as the rate in the presence of 161 mM and the H2S(g) reactive absorption efficiency can still be quite high, according to the results of Gendel (2007). The synergistic effect of Fe(III) and Cu(II) is discussed in the following Subsection. 4.2.6.3. Effect of initial concentration of Cu(II) As mentioned in Section 1.3.3 the catalytic effect of Cu(II) on Fe(II) oxidation is well known. Cher and Davidson (1955) observed that the catalytic effect in phosphoric acid solutions was significant for Cu2+ concentration up to 10−5 M, but there was little change with increasing the Cu2+ concentration above 10−3 M. This limited catalysis by Cu2+ was also observed by Zhang et al. (2000), who showed that the rate of Fe2+ 69 oxidation is significantly enhanced by adding 0.02 M CuSO4, but a further increase in the Cu2+ concentration to 0.1 M results in no additional change (Fig. 20). Figure 20. Effect of Cu2+ on Fe2+ oxidation by O2 (0.1 M FeSO4; pH=1.5; 80°C). (Source: Zhang et al., 2000) Similar results were obtained in the current work. In Fig. 21 the oxidation rates of three oxidation experiment series (each series with a different initial concentrations of Fe(III)) are presented. It is noticeable that the Fe(II) oxidation rate increased with an increase in Cu(II) concentration, but above ~4mM of Cu(II) and in the presence of Fe(III) the accelerating effect leveled off. This catalytic action is believed to be initiated by the reaction given in (49a) (Cher and Davidson, 1955; Zhang et al., 2000): a ,b Fe 2+ + Cu 2+ ←→ Fe 3+ + Cu + (49a,b) and followed by the reactions: Cu + + O2 → CuO2 + + CuO 2 + H + → Cu 2+ + HO2 (50) • (51) It was assumed that the first reaction (Eq. (49a)) is the rate determining step and that HO2• further reacts with Fe(II) as in the Weiss mechanism (Huffman and Davidson, 1956). It was concluded that the catalytic action of copper ion on the oxidation of Fe2+ by O2 relies on the fact that Cu+ reacts with O2 more readily than Fe2+ to generate HO2• and H2O2, and that the oxidation of Fe2+ ion by H2O2 is faster than by O2. 70 300 -d[Fe(II)]/dt [mg Fe/l/h] [Fe(III)]=0 250 a [Fe(III)]=81 mM [Fe(III)]=161 mM 200 150 100 50 0 0 2 4 6 8 10 12 14 16 Cu(II) concentration [mM] -d[Fe(II)]/dt [mg Fe/l/h] 80 b 60 40 [Fe(III)]=0 20 [Fe(III)]=81 mM [Fe(III)]=161 mM 0 0 2 4 6 8 10 12 Cu(II) concentration [mM] 14 16 Figure 21. Fe(II) oxidation rate as a function of the initial Cu(II) concentration with various concentrations of Fe(III). The × signs represent oxidation experiments that were conducted with pure oxygen (as opposed to oxidation with air). PT=581mM However, according to Zhang et al. (2000) Cu(I) further reacts with HO2• to form H2O2 (Eq. (52a)): • a ,b Cu + + HO2 + H + ←→ Cu 2+ + H 2 O2 (52) This phenomenon is apparently due to competing oxidation of Cu2+ according to Equations (52b) and (53). • HO 2 + Cu 2+ → O2 + Cu + + H + 71 (53) Zhang et al. (2000) also found that a high Cu2+ concentration (five times higher than the initial Fe2+ concentration) does accelerate the initial rate of Fe2+ oxidation, but has little effect once significant Fe3+ has been produced in the reaction mixture. The explanation for this is as follows: it appears that at relatively high initial concentrations of both Cu2+ and Fe2+, the oxidation of Fe2+ occurs via the generation of Cu+ (Eq. (49a)), and the subsequent production of H2O2. However, as the oxidation proceeds, Fe3+ builds up and the reverse reaction between Fe3+ and Cu+ (Eq. (49b)) becomes predominant. Thus, Fe3+ competes with O2 to oxidize Cu+ and as a result, lessens or disables the formation of CuO2+ that is the key intermediate of the proposed mechanism. In order to evaluate the effect of O2 on Cu(II) catalysis, three experiments were conducted with pure oxygen. The solutions contained 0.8 mM Cu(II) and the initial Fe(III) concentrations were 0, 0.08 and 0.16 M (represented as × signs in Fig 21). In the absence of Fe(III) the oxidation rate was similar to the oxidation rate attained with air (see Fig. 21b), which means that the oxidation rate is not limited by the concentration of O2. However, in the presence of Fe(III) the oxidation rate was higher than the oxidation rate with air (Fig. 21b). This result seems to indicate that the increase in O2 concentration increased the oxidation rate of Cu+ by O2 rather then the competing reaction with Fe3+, which resulted in a faster total-oxidation rate of Fe(II). Cher and Davidson (1955) developed the following rate equation according to their findings in phosphoric acid solution: [ ][ ] [ ] PO2 Fe 2+ Cu 2+ PO2 k1 Fe 2+ PO2 1 + k1 + k 3 Fe 3+ Fe 3+ d [Fe( II )] − = PO2 dt Cu 2+ 1 + k2 + k 4 Fe 3+ Fe 2+ [ ] [ ] [ ] [ [ ] ] (54) The saturation effect with respect to Cu2+ and the retarding effect of Fe3+ were observed also in sulfuric acid solutions (Lowson, 1982). However, the rate constant for the first step in the cupric ion catalyzed reaction (Eq. (49a)) in phosphoric acid is two-orders of magnitude higher than in sulfuric acid solution (Huffman and Davidson, 1956). The catalytic effect of Cu2+ in acidic ferrous chloride solutions had also been noted (Lowson, 1982). To summarize, the catalytic effect of Cu(II) is limited in the presence of Fe(III), and in the range of conditions that was examined in the current study, increasing the Cu(II) concentration above 4 mM results in an insignificant increase in the rate of Fe(II) 72 oxidation. The use of pure oxygen acts to increase the Fe(II) oxidation rate, however it is not cost effective and cannot be considered a process option.. 4.2.6.4. Effect of total phosphate concentration As discussed in Section 1.3.3 phosphate was found to increase the Fe(II) oxidation rate due to its complexing effect with the Fe(III) product (Cher and Davidson, 1955; King and Davidson, 1958). In Fig. 22 the oxidation rates of seven oxidation experiment series (each series with different initial concentrations of Fe(III) and Cu(II)) are presented. It is clear from the results that phosphate has a catalytic effect on the Fe(II) oxidation rate. A second order dependency on the total phosphate concentration (PT) was found to fit well the results, irrespective of the Fe(III) and Cu(II) concentrations (0.926<R2<1.000). In order to better understand the results, an analysis of the aqueous phase distribution of phosphate species in solution with respect to PT was made with the software MINEQL+. The sets of the input data for this theoretical analysis are given in Table 11. The distribution of the Fe(III) species and Fe(II) species for set D as a function of the PT are presented in Figures 23 and 24 respectively. Table 11. Composition of solutions that were analyzed by the MINEQL+ software for the distribution of phosphate species Concentrations [mM] Fe(II) Fe(III) Cu(II) H2SO4* Set A 18 0 0 100 to 35 Set B 18 0 8 75 to 45 Set C 18 161 0 400 to 200 Set D 18 161 8 400 to 300 * At a higher H3PO4 concentration in solution less H2SO4 was needed to reach pH1.0 73 80 [Fe(III)]=0, [Cu]=0 [Fe(III)]=0, [Cu]=7.9 [Fe(III)]=81, [Cu]=0.8 [Fe(III)]=81, [Cu]=1.6 [Fe(III)]=81, [Cu]=7.9 [Fe(III)]=161, [Cu]=0.8 [Fe(III)]=322, [Cu]=7.9 -d[Fe(II)]/dt [mg Fe/l/h] 70 60 50 in mM 40 30 20 10 0 0 100 200 300 400 500 600 700 Total phosphate concentration [mM] 800 900 Figure 22. Fe(II) oxidation rate as a function of PT with various concentrations of Fe(III) and Cu(II) Figure 23. Distribution of the Fe(III) species as a function of PT, in the presence of 1 g Fe(II)/l (18 mM), 9 g Fe(III)/l (161 mM) and 0.5 g Cu(II)/l (8 mM). pH1.0. Semilogarithmic scale 74 Figure 24. Distribution of the Fe(II) species as a function of PT, in the presence of 1 g Fe(II)/l (18 mM), 9 g Fe(III)/l (161 mM) and 0.5 g Cu(II)/l (8 mM). pH1.0. Semilogarithmic scale The analysis show that the dominant species of Fe(III) in the presence of phosphate is FeH2PO42+ (at PT > 0.12 M), and the dominant species of Fe(II) is FeH2PO4+ (at PT > 0.18 M). In addition second-order dependence was found to fit well the increase in concentrations of the species FeH2PO4+ and the ion H2PO4- with respect to PT (R2=0.9986 and 0.9982 respectively). The tendency was similar in the other simulations as well (sets A to C). Cher and Davidson (1955) and King and Davidson (1958) demonstrated that the oxidation of ferrous sulfate by molecular oxygen in phosphoric acid and pyrophosphoric acid solutions (at pH ~1-2, with the ionic strength adjusted to 1.0-1.1 with sodium perchlorate, and at 30°C) was second order with respect to H2PO4- and first order with respect to H2P2O72-, yielding the following overall rate law − d [Fe( II )] − = k1 Fe 2+ PO2 H 2 PO4 dt [ ] [ ] + k [Fe ]P [H P O ] 2 2− 2+ 2 O2 2 2 7 (55) King and Davidson (1958) concluded that the second-order dependence on phosphate was not due to the equilibrium − 2 H 2 PO4 ↔ H 2 P2 O7 2− + H 2O (56) 75 Thus, the second order dependency with respect to H2PO4- concentration, which was found in the current work, is corroborated by literature. No inhibition by Fe(III) was observed by Cher and Davidson (1955), and it was suggested that the complexing effect of Fe(III) by phosphate decreased the rate of Fe(III) reduction by HO2• (Eq. 46b). However, there was no Fe(III) at the beginning of the experiment of Cher and Davidson (1955), while in current work the initial concentration of Fe(III) was high and as a result inhibition was observed right from the start of the experiments. All in all, the results show a positive effect of phosphate on the Fe(II) oxidation rate, probably related to the concentration of the H2PO4- ion. The drawback of using a high PT is its rapid precipitation with Fe(III) even at the low pH value practiced in the current work (see Section 4.1.2.). 4.2.6.5. Effect of total sulfate concentration The oxidation of ferrous sulfate in sulfuric acid solution has been described as follows (Chmielewski and Charewicz, 1984; Rönnholm et al., 1999): 4 FeSO4 + O2 + 2 HSO4 + 2 H 3 O + → 2 Fe2 (SO4 )3 + 4 H 2 O − (57) Most of the previous studies indicate that at ambient temperatures and in sulfuric acid solutions below pH2, the rate of reaction is very slow, independent of pH, first order with respect to the partial pressure of oxygen and second order with respect to ferrous. Combined, these results are summarized in the following general rate equation (Lowson, 1982; Chmielewski and Charewicz, 1984): − 2 d [Fe( II )] = k Fe 2+ PO2 dt [ ] (58) Huffman and Davidson (1956) observed an increase in the reaction rate with sulfate concentration, at a fixed pH and ionic strength. .The rate was found to be independent of dissolved ferric ion provided that additional sulfuric acid is added; otherwise the product slows down the reaction rate (Macejevskis and Liepina, 1965, quoted by Lowson 1982). The presence of sulfate ions at near-neutral pH (as discussed in 1.3.2) decreased the oxidation rate, presumably because of the formation of the FeSO40 complex, which does not oxidize rapidly as compared to the complexes with hydroxyls and carbonates. It is possible that at low pH levels, where the concentrations of hydroxyls and carbonates are negligible, the oxidation of FeSO40 76 becomes significant. This hypothesis is supported by Dreisinger and Peters (1989) who reported on higher reactivity of the ferrous sulfate ion pair relative to the free ferrous ion, and Willix, (1963) and Ciavatta et al., (2002), who found that at low pH the product FeSO4+ is stable. Huffman and Davidson (1956) reported that the oxidation is accomplished by both catalyzed and un-catalyzed reactions, as described in the following rate expression: − 2 d (Fe( II ) ) 2 = k u Fe 2+ PO2 + k s (FeSO4 ) Fe 2+ PO2 = k t (Fe( II ) ) PO2 dt ( ) ( ) (59) In the current study the total sulfate concentration (ST) in the tested solutions was between 0.04 and 0.53 M (1.1 to 16.9 g S/l), resulting from the concentrations of the reagents added ((NH4)2Fe(SO4)2•6H2O, CuSO4, Fe2(SO4)3 and H2SO4). Fe(II) oxidation rates as a function of ST are presented in Fig. 25. The results are quite scattered, thus it can be only concluded that sulfate has a minor effect on the oxidation rate of Fe(II). The somewhat decreased reaction rate observed with an increase in ST can be attributed to the increase in Fe(III) concentration, that originates from the reagent Fe2(SO4)3. 300 in mM P=0, Cu=0 P=0, Cu=7.9 P=452, Fe(III)=81 P=581, Fe(III)=0 P=581, Fe(III)=81 P=581, Fe(III)=161 -d[Fe(II)]/dt [mg Fe/l/h] 250 200 150 100 50 0 0 50 100 150 200 250 300 350 400 450 Total sulfate concentration [mM] Figure 25. Fe(II) oxidation rate as a function of ST with various concentrations Fe(III), Cu(II) and PT 77 Summary: In the investigation of the catalytic oxidation of Fe(II) by oxygen the following results were obtained: The Fe(II) oxidation rate decreases with a decrease in Fe(III) concentration. This indicates that a low Fe(III) concentration is requested in order to enhance Fe(II) oxidation kinetics. However, a high concentration of Fe(III) is needed for efficient reactive-absorption of H2S(g), and thus this option is not practical. The Fe(II) oxidation rate is accelerated in the presence of Cu(II). The use of pure oxygen rather than air may further increase the oxidation rate. However, H2S reacts with Cu(II) and precipitate as CuS. Under the investigated condition CuS did not oxidize to S0 by Fe(III) as reported by Zhang and Tong (2006). Moreover, the Fe(II) oxidation rate is accelerated in the presence of phosphate. Yet, the phosphate concentration that is needed in order to achieve a reasonable oxidation rate (higher or equal to the oxidation rate of H2S by Fe(III)) is not feasible, since massive precipitation of P-Fe(III) species will occur. At pH1.0, in the presence of 4.5-9.0 g Fe(III)/l (81-161 mM) and less than 14 g P/l (452 mM) the maximal oxidation rate achieved was 7.5 mg Fe(II)/h. With this rate it is possible to treat up to 2.3 mg H2S/h, as calculated according to Eq. (60). rH 2 S( g ) g H 2S g Fe ( II ) rFe ( II ) × M .W . H 2 S g H 2 S (60) g Fe ( II ) gmol H 2 S h = = rFe ( II ) × 0.30515 g Fe ( II ) gmol Fe ( II ) h g Fe ( II ) M .W . Fe ( II ) ×E × 2 gmol Fe ( II ) gmol H 2 S where rH 2 S( g ) is the rate of H2S(aq) oxidation, rFe ( II ) is the rate of Fe(II) oxidation and E is the reactive-absorption efficiency of H2S(g) (assumed to be 1). All in all, the oxidation of Fe(II) by O2 at pH1.0 can be catalyzed by phosphoric acid and copper ions. However, this procedure cannot be considered feasible for the process of H2S(g) removal. 78 4.3. Electrochemical Fe(II) oxidation as part of the LRSR process 4.3.1. H2S(g) reactive-absorption efficiency A PVC pipe was used as a bubble column reactor (internal diameter 158 mm). A sintered-glass plate diffuser (5 inches diameter) with a pore size distribution of 150 200 µm was used to generate the bubbles. The effect of pH and chloride concentration was tested with an air flux of 1.1 m3air/min/m3solution and the results are presented in Fig. 26 and Fig. 27 as a function of the inlet H2S(g) concentration. It seems that both increase in pH and increase in the chloride concentration in solution have a positive effect on the efficiency of H2S(g) reactive-absorption. However, the effect of chloride was much stronger in the range of conditions tested: while increasing the pH from 1.0 to 1.7 elevated the reactive-absorption efficiency from 80% to 89% in average, the addition of 0.28 M Cl- elevated the efficiency from 18% to 70% in average. H2S(g) reactive-absorption efficiency [%] 95 pH=1.74 pH=1.38 pH=1.00 90 85 80 75 0 10 20 30 40 H2S(g) concentration [ppm] 50 60 Figure 26. Effect of pH on the H2S(g) reactive-absorption efficiency. [Cl-] = 0.85M, air flux = 1.1 m3/m3/min 79 H2S(g) reactive-absorption efficiency [%] 90 80 70 0 M Cl0.28 M Cl0.85 M Cl1.41 M Cl- 60 50 40 30 20 10 0 0 10 20 30 40 H2S(g) concentraion [ppm] 50 60 Figure 27. Effect of Cl- concentration on the H2S(g) reactive-absorption efficiency. pH1.0, air flux = 1.1 m3/m3/min Fig. 28 presents the distribution of the Fe(III) species as a function of pH in the presence of 0.85 M Cl- (30 g Cl/l) (analysis using MINEQL+). The concentrations of the species FeOH2+, Fe(OH)2+ and Fe2(OH)24+ increase with an increase in pH. The concentration of the species Fe3(OH)45+ also increases significantly, from 10-9.1 to 106.4 M. The same tendency was found also in solution in the absence of Cl-. The species FeOH2+ was reported by Asai et al. (1990) and Ebrahimi et al. (2003) to oxidize H2S(aq). Ebrahimi et al. (2003) hypothesized that the species Fe3(OH)45+ and Fe2(OH)24+ also oxidize H2S(aq) (see Subsection 1.2.4.2). Thus it can be safely concluded that the increase in the reactive-absorption efficiency with pH is a result of the increase in the concentrations of the Fe(III) hydroxyl complexes. The effect of chloride concentration on the H2S(g) reactive-absorption efficiency was also analyzed by MINEQL+ while pH was maintained constant. The results are presented in Fig. 29. 80 Figure 28. Distribution of Fe(III) species as a function of pH in the presence of Cl- (0.85 M). Semi-logarithmic scale. Figure 29a. Distribution of Fe(III) species as a function of chloride concentration. Semi-logarithmic scale. pH1.0 81 H2S(g) reactive-absorption efficiency [%] Figure 29b. Distribution of Fe(III) species as a function of chloride concentration. Normal scale. pH1.0 According to Fig. 29a the concentrations of the hydroxyl complexes, which were found to oxidize H2S(aq), decrease with the increase in total-Cl- concentration. It could have been assumed that the H2S(g) reactive-absorption efficiency would decrease with the increase in Cl- concentration. However, the reactive-absorption efficiency of H2S(g) in the presence of Cl- was found to be much higher than the efficiency in the absence of Cl- (Fig.s 27 and 29b). Even at low total-[Cl-], the concentrations of the Fe3+Cl- complexes are high and increase with an increase in total-[Cl-]. It is thus concluded that the Fe3+Cl- complexes can also oxidize well H2S(aq). No information can be found in the literature on reactive-absorption of H2S(g) in the presence of chlorides, and the observation that the Fe3+Cl- complexes can oxidize H2S(aq) is new. In Fig. 29b the average values of H2S(g) reactive-absorption efficiencies, for each one of the experiments presented in Fig. 27, are also presented. It is noticeable that the dependency of the reactive-absorption efficiency on the total-[Cl-] is similar to the dependency of the concentration of the FeCl2+ species on total-[Cl-]. It is thus not unlikely that the dominant oxidizing complex of H2S(g) at pH1.0 is FeCl2+. The main conclusion from this section is that the use of a solution that does not contain chlorides is not applicable for the LRSR process at pH1.0. A second 82 conclusion is that increasing the Cl- concentration beyond 30 g/l (0.85 M) does not improve the reactive-absorption efficiency of H2S(aq) and seems thus to be superfluous. 4.3.2. Practical precipitation potential in the working solution The effect of pH and chloride concentration was tested. The compositions of the tested solutions are listed in Table 12. Table 12. Conditions of solutions for precipitation experiments (electrochemical oxidation) and mass of precipitates after 152 days Bottles No. Concentrations [mM] Fe(II) A - Fe(III) Cl 161 846 (=9.0 g Fe/l) (=30 g Fe/l) 161 pH Mass of precipitate [mg per l-solution] 1.0 8 846 1.4 56 161 846 1.7 528 161 282 1.0 31 E 161 282 1.4 34 F 81 846 1.0 45 B C D 18 (=1 g Fe/l) From a thermodynamic standpoint and based on MINEQL+ simulations all the considered solutions were projected to precipitate Fe(III) solids to a degree that the soluble FeT concentration remaining in solution at equilibrium would be very low. For example, in all the experiments which were conducted at pH1.0 (A, D and F) more than 94% of the Fe(III) was supposed to precipitate as part of the solid Fe(OH)2.7Cl0.3. In practice, however, the pH value was found the most dominant parameter in the formation of ferric precipitates under the investigated conditions. After a period of 51 days all solutions seemed visually clear. However, on the walls of bottle C (0.85 M Cl- at pH1.7) a white precipitate could be observed (Fig. 30). No change in pH was detected. After 138 days from the beginning of the experiment, precipitation has increased in bottle C and small suspended clusters could be seen in solution (although no turbidity was observed). In bottle B (0.85 M Cl- at pH1.4) very 83 slight precipitation was observed. After 152 days the solutions (<500 ml) were filtered through a 0.45 µm filter paper and the filters were dried in 105°C for 24 hours. The mass of precipitates that were found are presented in Table 12. After 152 days a significant amount of precipitates (528 mg per l-solution) were observed in bottle C. In the bottles that were held at pH1.0 and pH1.4 the amount of precipitates was much lower (more than one order of magnitude). Yet, at pH1.4 the amount of precipitates was higher than at pH pH1.0. It can be concluded that although precipitation was anticipated by thermodynamic data, the kinetics of precipitation was low and affected by pH. The pH of solution should not be higher than pH1.4 for a long-term LRSR process. Working at pH1.0 provides a safety factor. C B Figure 30. Bottles from precipitation experiments after 51 days. B (pH1.4) and C (pH1.7). [Cl-]=0.85 M 4.3.3. Fe(II) oxidation rate in the electrochemical oxidation process The majority of the experiments were performed using a constant current (±0.1 amp) and the voltage values were measured. Otherwise, voltage was kept constant and the current was measured. All the electro-oxidation experiments were conducted at pH1.0. In earlier experiments (Subsections 4.2.3.1 and 4.2.3.2) the cathode was made of titanium coated with ruthenium oxide (Ti/RuO2). Later on it was decided to replace the electrode to bare Ti, since the RuO2 coating started to detach from the electrode after a few hours of operation. A bare Ti electrode was used in the experiments shown in Subsections 4.2.3.3 and 4.2.3.4. 84 4.3.3.1. Direct electrooxidation The rate of direct oxidation was evaluated and compared to the rate of indirect oxidation. The conditions of the experiments are listed in Table 13 and the results are shown in Fig. 31. The total iron concentration in the experiments was 0.179 M (10 g Fe/l). Although the Fe(III) concentration in the proposed process are designed to vary inside the 9-8 g Fe/l interval, initial Fe(II) and Fe(III) concentrations were both set at 5 g Fe/l (0.090 M) in order to obtain the influence of the Fe(III) concentration on the Fe(II) electro-oxidation rate. It is noticeable from Fig. 31 that the oxidation rate of Fe(II) is higher at lower [Fe(III)]:[Fe(II)] ratio. In fact, there were two tries to perform direct oxidation in solution with an initial [Fe(III)]:[Fe(II)] ratio of 9:1. These experiments were stopped after a short period (about 30 minutes) since the concentration of Fe(II) increased rather than decreased with time. It was concluded that Fe(III), being at such high concentration, competes with H+ on reduction on the cathode surface and thus under such conditions the back reduction reaction of Fe(III) to Fe(II) was faster than the forward reaction (Fe(II) → Fe(III)). Yet, high ratio of [Fe(III)]:[Fe(II)] is essential in order to achieve efficient reactive-absorption of H2S(g) (see 1.2.4.2). The Fe(II) oxidation rate recorded for the direct oxidation was significantly lower than the rates attained in the indirect oxidation experiments: 0.22 g/h in first two hours for direct oxidation compared to 1.41 g/h for indirect oxidation. Nevertheless, these rates are much higher than the rates that were obtained in the catalytic oxidation (Section 4.1.3). Since the preliminary experiments of indirect oxidation showed higher oxidation rates than direct oxidation and since the reactive-absorption efficiency of H2S(g) was very low in the absence of chlorides (Section 4.2.1), the indirect oxidation option seemed, from the inception, much more promising as the regenerative step for Fe(III) within the LRSR process at pH1.0. 4.3.3.2. Indirect electrooxidation - effect of anode to cathode surface area ratio (Sa:Sc) High anode to cathode surface area ratio (Sa:Sc) may suppress the Fe(III) back reduction on the cathode, which is important because of the need for high Fe(III) 85 concentration in the LRSR solution. On the other hand, high Sa:Sc may suppress the wanted H2 evolution on the cathode. Two surface area ratios (Sa:Sc) were examined: 7.1 (smaller cathode) and 1.7 (bigger cathode). The conditions of these experiments are listed in Table 13 and the results are shown in Fig. 31. It is very clear from Fig. 31 that the net Fe(II) oxidation rate was higher in the experiments with lower Sa:Sc ratio (bigger cathode). This means that H2 evolution was more dominant than Fe(III) back reduction on the cathode surface, and increasing the cathode surface allowed higher H2 evolution rate and consequently higher Fe(II) oxidation rate. This is supported by Alvarez-Gallegos and Pletcher (1998) and Bisang (2000), who reported that the main reaction on the cathode around pH2 was H2 evolution (see Subsection 1.4.2.1). Table 13. Conditions of electrochemical oxidation experiments conducted with a Ti/RuO2 anode Direct / indirect oxidation Direct Indirect Indirect Indirect Indirect Initial Current Voltage pH [amp] [V] 1.01 0.3 2a 0.95 1.03 Cl- Cathode concentration dimensions [M] [cm2] b 0 : 10 0 0.72 × 9.15 Fe(III):Fe(II) a 1 3.5 5:5 0.85 0.72 × 9.15 1 3 a 5:5 0.85 0.72 × 9.15 3 a 5:5 0.85 3.0 × 9.14 2.6 5:5 0.85 3.0 × 9.14 1.0 2.2 1.0 a 2 a was kept constant during the experiment b anode dimensions 5.1 x 9.15 cm2 in all experiments 86 Fe(II) concentration [g Fe/l] Figure 31. Concentration of Fe(II) as a function of time in the experiments conducted with Ti/RuO2 cathode 4.3.3.3. Indirect electrooxidation - effect of current density The range of currents that were investigated in literature is between 0.2 and 4.2 kA/m2 (Robertson et al., 1983; Kelsall, 1984; Rudolf et al., 1995; Khelifa et al., 2004). An increase in the active chlorine production rate as a result of increase in the current density was reported by Rudolf et al. (1995), Kraft et al., (1999) and Khelifa et al. (2004). The effect of current density in the current work was tested with a Ti cathode and Ti/RuO2 anode, both with dimensions of 9.15×5.1 cm2. The range of current densities was between 0.43 and 1.71 kA/m2 (2 to 8 A with 9.15×5.1 cm2 electrodes). The chloride concentration in all the experiments was 0.85 M. This concentration was chosen according to the results of the H2S(g) reactive-absorption efficiency experiments (Section 4.2.1). The interesting interval of Fe(II) concentration is between 2 to 1 g Fe/l from the standpoint of the proposed LRSR process (and Fe(III) concentration is 8-9 g/l). Thus, the Fe(II) oxidation rate was calculated in this interval by a linear regression. For example, the incline of the linear regression for the experiment conducted with 0.43 kA/m2 was -1.1356 g Fe(II)/l/h (Fig. 32). The solution volume was 1.5 l, thus the Fe(II) oxidation rate was 1.70 g Fe(II)/h, as seen in Fig. 33. 87 5 0.43 kA/m2 0.54 Fe(II) concentration [g Fe/l] 4 kA/m2 0.64 kA/m2 0.86 kA/m2 1.07 kA/m2 1.29 kA/m2 1.71 kA/m2 3 2 y = -1.1356x + 3.9024 R2 = 0.9946 1 0 0.0 0.5 1.0 1.5 2.0 Time [h] 2.5 3.0 3.5 Figure 32. Concentration of Fe(II) as a function of time in the indirect 72 7 64 6 56 48 40 4 32 3 24 2 -d[Fe(II)]/dt Voltage 1 0 0.0 16 Current efficiency [%] 8 5 Voltage [V] -d[Fe(II)]/dt [g Fe/h] electrochemical experiments at different current densities. [Cl-]=0.85 M Current efficiency 0.5 1.0 1.5 2 Current density [kA/m ] 8 0 2.0 Figure 33. Effect of current density on Fe(II) oxidation rate, voltage and current efficiency. [Cl-]=0.85 M 88 Fig. 33 shows the Fe(II) oxidation rates (in the interesting interval) and the current efficiencies as a function of the current densities. Higher Fe(II) oxidation rates were achieved for higher current densities. The highest Fe(II) oxidation rate recorded was 7.1 gFe/h, and it corresponded to current efficiency of 42.8%. The highest current efficiency (59%) was achieved at 0.64 kA/m2. 4.3.3.4. Indirect electrooxidation - effect of chloride concentration Khelifa et al. (2004) examined the effect of chlorine concentration in the range of 6 to 150 g-Cl/l and reported on an increase in active chlorine production with an increase in chloride concentration. The same tendency was observed by Krstajić et al. (1991) who examined the hypochlorite production rate for NaCl solutions with concentrations of 14.6 g/l, 19.28 g/l and 29.21 g/l. However, it was also reported by Khelifa et al. (2004) that at NaCl concentration above 100 g-Cl/l the active chlorine production rate levels off. Rudolf et al. (1995) reported that higher NaCl concentrations decreased chlorate formation in divided electrolytic cell. The same was observed by Czarnetzki and Janssen (1992) who performed their experiments in an electrolytic cell divided by anionic membrane. The effect of chloride concentration was tested with a Ti cathode and Ti/RuO2 anode, both with dimensions of 9.15×5.1 cm2. The range of chloride concentrations in these experiments was between 0.28 and 2.26 kA/m2 (10-80 g Cl/l). All experiments were conducted with current density of 0.64 kA/m2, which was shown to result in the highest current efficiency in the presence of 0.85 M Cl- (Subsection 4.2.3.3). The oxidation rates were determined by linear regression of the –d[Fe(II)]/dt curve in the interval of 2 to1 g Fe(II)/l. The results are presented in Fig. 34 and 35. The voltage measured in the experiments decreased with the Cl- concentration, which means that for higher Cl- concentration lower voltage is required to obtain the same current density. The increase in chloride ions concentration from 0.28 to 0.56 M increased significantly both the oxidation rate and the current efficiency. At Cl- concentrations higher than 0.85 M the effect was minor, similar to the results reported by Khelifa et al. (2004). It is thus plausible to assume that at high Cl- concentrations the Fe(II) concentration becomes the limiting factor of the Fe(II) oxidation reaction. 89 Fe(II) concentration [g Fe/l] Figure 34. Concentration of Fe(II) as a function of time in the indirect electrochemical 4.5 72 4.0 64 3.5 56 3.0 48 2.5 40 2.0 32 1.5 24 -d[Fe(II)]/dt voltage 1.0 16 Current efficiency 0.5 Current efficiency [%] Voltage [V] -d[Fe(II)]/dt [g Fe/h] experiments with different chloride concentrations. Idensity=0.64 kA/m2 8 0.0 0 0.0 0.5 1.0 1.5 2.0 Concentration of Cl- [M] 2.5 Figure 35. Effect of Cl- concentration on Fe(II) oxidation rate, voltage and current efficiency 90 4.3.3.5. Potential chlorine loss Chlorine escape during the electrolysis was measured in the experiments that were conducted with varied current densities and Cl- concentraion of 0.85 M (Subsection 4.2.3.3). Chlorine escape was observed only in the experiments with current densities of 0.64 and 1.71 kA/m2. The amounts of Cl2(g) lost were 0.2 and 0.4 mM respectively, less than 0.06% of total Cl- concentration. It should also be noted that both experiments lasted much longer the point where all the Fe(II) was oxidized. Industrial scale reactors should however take this phenomenon into consideration. 4.3.3.6. Indirect electrooxidation - energy cost The highest current efficiency (59%) was found at current density of 0.64 kA/m2 in the presence of 30 g Cl-/l (0.85 M). The voltage was 2.87 V and the oxidation rate was 3.7 g Fe(II)/h, which is equivalent to 1.13 g H2S/h (three orders of magnitude higher than the rate obtained in the catalytic oxidation). The energy needed under such conditions is 7.6×10-3 kW·h normalized for the removal of 1 gram of H2S(g), calculated according to Eq. (61). ( 2.87 V × 0.64 kA ( )) × 9.15 × 5.1 cm 2 kW ⋅ h m2 = 7.6 * 10 −3 gH S g H2S 1.13 2 h (61) Assuming that the cost of electricity is 0.12 $/kW·h the cost of removing 1 kg of H2S(g) would be 0.91 $, which does not appear excessive. 7.6 * 10 − 3 kW ⋅ h $ $ $ × 0.12 = 9.1 * 10 − 4 = 0.91 g H 2S kW ⋅ h g H 2S kg H 2 S (62) 1kg of H2S corresponds, for example, to 1455 m3 of air contaminated with 500 ppm of H2S (a relatively high concentration). mg H 2 S 500 ppm × 1.375 m 3 air × 1455m 3 = 1kg air H 2S ppm H 2 S (63) 4.3.3.7. Volume of solution in the electrolytic reactor An electrolytic cell that is designed to oxidize, for example, 3.7 g Fe(II)/h, can remove 1.13 g H2S/h. The amount of 1.13 g H2S corresponds to 1636 l of air contaminated with 500 ppm of H2S, which can be treated by a recycle flow rate of 27.3 l/min. If the flux of air is set at 1.1 m3air/min/m3solution the volume of the 91 absorbing solution should be about 25 l. The volume of the electrolytic cell used in the experiments reported in the current work was about 60 ml. 1.13 g H 2S h × 1455 l air g H 2S 1.1l air / min/ l solution 1636 = l air h l air min = 24.8l solution (64) / min/ l solution 27.3 1.1l air / min/ l solution × 60 min = h 1.1l air 4.3.3.8. Electrode material Titanium, which was used as the cathode in the electrochemical oxidation experiments, is known to have a high corrosion resistance in active chlorine solutions. At acidic conditions (pH<5.0), when cathodic polarization is applied, the formation of titanium hydride (TiH2) was reported to occur (Videm et al., 2008), which leads to the embrittlement of titanium. Consequently the use of titanium as hydrogen evolution cathode under acidic conditions is restricted in industrial long term operations. An alternative cathode material was not attempted. 4.3.3.9. Changes in pH An increase in pH from pH1.0 to pH1.2 was observed in most experiments, due to H2 evolution on the cathode. Rudolf et al. (1995) observed that the pH of an electrolyzed NaCl solution increased during electrolysis from an initial pH6 to pH9.0-9.2 after several minutes finally stabilizing at the range pH9.55±0.15 after another 20 minutes. From the standpoint of the total process this increase in pH should not be a problem, since the oxidation of H2S to S0 increases H+ concentration in the solution at an equal rate. To summarize all the above findings, indirect electro-oxidation of Fe(II) was found much more suitable than direct oxidation for the H2S(g) removal process at pH1.0, due to higher Fe(II) oxidation rates and significantly higher efficiency of H2S(g) reactiveabsorption in the presence of a high chloride concentration. The use of 30 g Cl-/l (0.85 M) seems to be the most advantageous chloride ion concentration, both from the electrochemical oxidation perspective and from the H2S(g) reactive-absorption standpoint. 92 10. Conclusions • Catalytic oxidation of Fe(II) by oxygen in the presence of copper and phosphate is possible but not applicable for an LRSR process operated at pH1.0, since the Cu(II) will precipitate with H2S and phosphate will precipitate with Fe(III). • The accelerating effect of phosphate on Fe(II) oxidation rate at pH1.0 is probably related to the concentration of the H2PO4- ion. • The accelerating effect of copper on the Fe(II) oxidation rate at pH1.0 increases in the presence of a higher O2 concentration and decreases in the presence of a higher Fe(III) concentration. • The presence of Fe(III) at high concentrations decreases the Fe(II) oxidation rate, both by O2 and electrochemically. • The reactive absorption efficiency of H2S(g) increases in the presence of chlorides, although the addition of more than 30 g Cl/l seems to be superfluous. It is possible that the positive effect of Cl- is due to high oxidation rate of H2S(aq) by the FeCl2+ complex. • Direct electrochemical oxidation of Fe(II) is not a viable option for an LRSR process at pH1.0, mainly because the reactive-absorption efficiency of H2S(g) is very low in the absence of chlorides. • Increase in current density (0.43 to 1.71 kA/m2) increases the rate of indirect electro-oxidation of Fe(II), although the highest current efficiency (59%) was found with current density of 0.64 kA/m2. • Increase in Cl- concentration up to 0.85 M (30 g Cl/l) increases the rate of indirect electro-oxidation of Fe(II). Above 0.85 M the oxidation rate becomes approximately constant. • Indirect electrochemical oxidation of Fe(II) is applicable for an LRSR process at pH1.0: the Fe(II) oxidation rates are very high, the H2S(g) reactive-absorption efficiency is high and precipitation is minimal even after 5 months. 93 Reference list Alvarez-Gallegos, A. and Pletcher, D. (1998). 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