amines is difficult by HPLC because of auto-oxidation of aromatic amines [47, 48]. The ...... In run 3, 11-th repeated culture was used for further three enrichment ...
Decolorization and Degradation of an Azo Dye, Orange II by Chemical and Biological Treatments
Department of Environmental Chemistry and Engineering Tokyo Institute of Technology Shizuo Nishide
Contents Chapter 1 Introduction
1.1. Synthetic dyes and the effluents ................................................................................... 1 1.2. The problem of azo dye waste ..................................................................................2 1.3. Physicochemical process ..........................................................................................3 1.3.1. Ozonation .........................................................................................................4 1.3.2. Fenton process ..................................................................................................4 1.3.3. Photocatalyst
(UV/TiO2) ...............................................................................5
1.4. Advanced biological treatemat .................................................................................6 1.4.1. Fungal degradation of dyes based on peroxidase .............................................7 1.4.2. Anaerobic decolorization by methanogenic granular sludge .........................8 1.4.3. Aerobic bacterial decolorization and degradation ..........................................9 1.5. The use of zero-valence metals for environmental contaminants ........................10 1.5.1. The reduction of azo dyes by zero valence iron ...........................................12 1.5.2. The combined use of zero-valence iron and biological treatment ...............13 1.6. Purpose of this study ............................................................................................14 References .....................................................................................................................17 Tables ............................................................................................................................25 Figures ...........................................................................................................................29
Chapter 2 Decolorization of azo dyes by PIP tin balls in citric acid solution 2.1. Introduction ..........................................................................................................32 2.2. Methods ................................................................................................................33
2.2.1. Materials ........................................................................................................33 2.2.2. Decolorization of azo dyes under static conditions ......................................34 2.2.3. HPLC analysis and product identification ...................................................34 2.2.4. Decolorization of OrangeⅡunder stirred condition .....................................35 2.2.5. Repeated decolorization of OrangeⅡunder anaerobic condition ................36 2.2.6. FT-IR measurement ......................................................................................37 2.2.7. XPS analysis .................................................................................................37 2.3. Results ..................................................................................................................37 2.3.1. Decolorization of azo dyes by PIP tin balls under static conditions ............37 2.3.2. Products after decolorization of Orange II ...................................................38 2.3.3. Effect of oxygen on the decolorization of Orange Ⅱ
................................39
2.3.4. Repeated decolorization of OrangeⅡ ...........................................................39 2.4. Discussion.............................................................................................................41 References ...................................................................................................................46 Tables ..........................................................................................................................48 Figures .........................................................................................................................53
Chapter 3 Comparative study on the decolorization of Orange II by citric acid and hydrochloric acid solutions 3.1. Introduction ..........................................................................................................62 3.2. Methods ................................................................................................................62 3.2.1. Materials .........................................................................................................62 3.2.2. Decolorization procedure by zero-valence tin .............................................63 3.2.3. Analysis ..........................................................................................................63
3.3. Results .....................................................................................................................64 3.3.1. Effect of acid concentrations on decolorization of Orange II ......................64 3.3.2. Effect of initial dye concentrations on Orange II decolorization .................66 3.3.3. Dependence of pH in citrate buffer on the decolorization of Orange II ......68 3.4. Discussion.............................................................................................................68 References ...................................................................................................................73 Tables ..........................................................................................................................75 Figures .........................................................................................................................79
Chapter 4 Biodegradation of two aromatic amines produced from Orange II by zero-valence tin 4.1. Introduction ..........................................................................................................90 4.2. Methods ................................................................................................................92 4.2.1. Materials .........................................................................................................91 4.2.2. Enrichment culture .......................................................................................91 4.2.3. Isolation of amine-degrading strains ............................................................92 4.2.4. Preservation of isolated strains .....................................................................92 4.2.5. Degradation of aromatic amines by isolated strains.......................................93 4.2.6. Preparation of tin-treated solution of Orange II .............................................93 4.2.7. Co-culture treatments of a mixture of authentic samples of ABS and 1A2N or tin-decolorized solution of Orange II .....................................94 4.2.8. Analysis ..........................................................................................................94 4.3. Results ..................................................................................................................95 4.3.1 Enrichment culture ........................................................................................95 4.3.2. Degradation kinetics of ABS by ABS-degrading consortium .....................98
4.3.3. Degradation of 1A2N ...................................................................................98 4.3.4 Co-culture of ABS-degrading consortium and m2 in the tin-treated solution of Orange II ........................................................99 4.4. Discussion...........................................................................................................101 References ............................................................................................................105 Tables ...................................................................................................................107 Figures ..................................................................................................................110
Chapter 5 Decolorization of Orange II by combined use of tin and Pseudomonas aeruginosa m2 in the presence of citrate 5.1. Introduction ........................................................................................................117 5.2. Methods ..............................................................................................................118 5.2.1. Preservation of P. aeruginosa m2 ................................................................ 118 5.2.2. Preculture with nutrient broth ...................................................................... 119 5.2.3. Mineral medium .......................................................................................... 119 5.2.4. Decolorization of Orange II by P. aeruginosa m2 ..................................... 119 5.2.5. Combined use of tin and m2 strain under shaking condition ....................120 5.2.6. Counting of the viable cell number .............................................................120 5.2.7. Analysis .......................................................................................................120 5.3. Results ................................................................................................................121 5.3.1. Decolorization of Orange II by m2 strain under static condition ..............121 5.3.2. Combined use of tin and m2 strain under shaking condition ....................122 5.4. Discussion...........................................................................................................124 References .................................................................................................................127 Tables .........................................................................................................................129
Figures .......................................................................................................................133
Chapter 6 Conclusions 6.1. Summary of results obtained from this study ...................................................138 6.1.1. Chapter 2 .....................................................................................................138 6.1.2. Chapter 3 .....................................................................................................138 6.1.3. Chapter 4 .....................................................................................................139 6.1.5. Chapter 5 .....................................................................................................140 6.2. Advantage of tin and biological methods developed in this study ...................140 6.3. Future aspects ...................................................................................................142 6.3.1. The recovery and recycle of tin .................................................................142 6.3.2. The application to the actual system............................................................142 6.3.3. The mechanism of the culture in the presence of tin .................................143 References .................................................................................................................144
Acknowledgements ...................................................................................................146
Chapter 1 1 Introduction
1.1. Synthetic dyes and the effluents Since the first synthetic dye, mauveine, was discovered by Perkin in 1856, thousands of synthetic dyes have been manufactured and synthetic dyes quickly replaced the traditional natural dyes. In 1860, the diazo coupling reaction of aromatic diazonium salts was discovered [1] and the resulting azo dyes, which were deep and bright color, were more than half of the annually produced amounts of dyes. Azo compounds have been widely used in textile, leather, cosmetic, plastic, food and recently in recordable optical disk such as CD-R, DVD-R, with the textile industry as the largest consumer. The usage of azo dyes will continue in the future. Since a small amount of dyes is clearly visible, the release of dye wastewaters in the environment is aesthetically unpleasant. In addition, many dyes are mutagenic and toxic and pose a serious hazard to aquatic living organisms. Colored wastewater is a consequence of processes both in the dye manufacturing industries and dye-consuming industries. Two per cents of dyes that are produced are discharged directly in aqueous effluent, and 10 % are subsequently lost during the textile coloration process. An indication of the scale of the problem is given by the observation that the annual market for dyes is more than 7×105 tonnes per year [2]. The degree of fixation for different dye and fiber combination and the loss to effluent are shown in Table 1-1[3]. Recently the popularity of cotton and the use of reactive dyes have increased. As reactive dyes have low rates of fixation, the dye effluent problem is deteriorated. Since environmental policy in the UK, in 1997, has stated that zero synthetic chemicals should be released 1
into the marine environment [2] in Europe, Government legislation with respect to colored effluent has been enforced. The trend is expanding into other countries. In Japan, the legislation was enforced by local governments in Wakayama city, Kawasaki city and Misato city. In the past, municipal sewage treatments were mainly used for the purification of textile mill wastewaters. These systems, however, depended mainly on biological activities and were found mostly inefficient in the removal of the more resistant synthetic dyes. As the combined effect of sewage treatment and dilution is not sufficient to remove the residual dye from process wastewater, and municipal treatment costs are increasing [4], an alternative industrial-scale decolorization system is needed to achieve the necessary color consent limits. Therefore, the color creating industries and color using industries are compelled to search for novel treatments and technologies that are directed particularly towards the decolorization of dyes in effluents.
1.2. The problem of azo dye waste Among several dyes, especially azo dyes can cause harmful effects to different organisms including humans. Most of the attention concerning the possible hazards arising from azo dye has been transferred to these reduction products, aromatic amines which were recognized as carcinogen and mutagen. Germany followed by the European Union published a list 22 aromatic amines classified as carcinogenic and banned the azo dyes that could form any of these amines by the cleavage of the azo bond. A regulation also came into force in Turkey, Netherland, India, France and China [5]. In Japan, a government regulation has not been enforced yet. However, Japanese textile industries have started self-imposed restraints on the products in relation with these defined
2
aromatic amines. Due to the general association of azo toxicity with hydrophobic character, low toxicity has often been achieved through the introduction of polar moieties in the dye structure, also resulting in higher aqueous solubility, sulfonation being widespread character [6]. However, this feature can hinder their removal in wastewater treatment works. Increased hydrophilicity has been as unfavorable for elimination of dye in activated sludge systems, through adsorption onto the biomass [7-9]. Furthermore, dye can be detrimental to the microbial population present in treatment works and may lead to decrease efficiency or treatment failure in plants [10, 11]. Conventional activated sludge systems were often combined with coagulation by alumina or ferric salt. However, the production of hazardous huge sludge was harmful [11, 12]. Addition of active carbon to activated sludge performed good color removal [13], but the drawback of activated carbon process is very high cost and the difficulty in regeneration [14]. More effective and more economical technologies than conventional biological methods are needed. Thus, physicochemical oxidation process, newly biological, and the combination of physicochemical and biological processes are proposed. The advantages and disadvantages of these methods are summarized in Table 1-2.
1.3. Physicochemical process Decolorization and mineralization of the persistent azo dyes using strong oxiding agents is considered to be rapid treatment process. Especially, ozonation, Fenton, TiO2/UV methods have been extensively studied.
3
1.3.1. Ozonation The use of ozone for purpose of wastewater treatment has some advantages. Ozone, which has high redox potential (2.07 V), can degrade many persistent organic pollutants such as chlorinated hydrocarbons, phenols, pesticides and aromatics [14, 15]. Ozone attacks conjugated double bonds and aromatic rings in the structure of dyes, resulting in the degradation of them into smaller molecules. Ozonation has no problem of second pollution, in which ozone converts to non-toxic oxygen through the reaction. One additional advantage is that ozone can be applied in its gaseous state and therefore, does not increase the volume of wastewater and sludge [14]. Ozonation is reported to be very effective for complete removal color and full plant works are gradually progressing in textile wastewater [9, 11, 16]. However, ozone is short half-life, typically being 20 min, and therefore continuous ozonation is required, which is major drawback in cost.
1.3.2.
Fenton process
The oxidation processes which are based on the generation of hydroxyl radicals are generally referred to advanced oxidation processes (AOPs). Hydroxyl radical has high oxidation potential (2.80 V) and can promote degradation of a wide range of contaminants in minutes. Hydroxyl radicals can be generated by UV/H2O2, O3/H2O2, O3/UV, TiO2/UV and Fe(II) salts/H2O2 (Fenton reagents) as shown in Fig. 1-1. Among these AOPs, Fenton process is comparatively cheap and easy to handle reagents [17]. Fenton reagent works well under acidic condition by oxidizing ferrous ion to ferric iron with simultaneous splitting of H2O2 into hydroxide ion and hydroxyl radical. Ferric ion returns to ferrous ion in the presence of H2O2, resulting in the chain reaction of the formation of hydroxyl radical.
4
Fenton process which is effective for color and COD removal of the dye effluent was reported that the cleavage of azo groups, conjugated unsaturated bond (-N=N-) in molecules was rapidly occurred [17,18]. Besides hydroxyl radical, the highly stable complex between Fe3+ and azo bonds, which leads to the photolysis in the presence of H2O2, was active in the cleavage of azo bonds.
The decolorization reaction was
followed by the degradation of aromatic ring and other moieties by hydroxyl radical [18]. The efficiency of Fenton reaction depends mainly on H2O2 concentration, Fe2+/H2O2 ratio, pH, and temperature. Under optimum conditions, the complete color removal was observed even in high strength wastewaters [19]. However, the major drawbacks are that flocculation by Fe3+ and dye molecules produces a large amount of hazardous sludge, and the removal of the aqueous reagent is required. To overcome these drawbacks, the heterogeneous Fenton catalytic degradation of azo dyes, which contained Fe ion immobilization, Fe clusters or Fe oxides, was studied to show efficient activity under neutral pH, which is favorable for the environment and low Fe leaching [20-23].
1.3.3. Photocatalyst (UV/TiO2) Major advantage of UV/TiO2 process is a semi-permanent use of the catalyst because titanium oxide is relatively high in chemical stability [24]. The process is initiated upon UV irradiation of the semiconductor with the formation of high energy electron/hole pairs by exciting an electron from the valence band (VB) to the conduction band (CB). The highly oxidative hole may directly react with the surface adsorbed organic molecules or indirectly oxidize the organic compounds via formation of hydroxyl
5
radical (Fig. 1). The recombination between hole and photogenarated electron is a major limiting factor. The couple of SnO2 or ZnO colloids with TiO2 is one interesting way to overcome the limitation. In the coupled semiconductor systems, photogenerated electron on TiO2 rapidly moves to the conduction band of other semiconductor so that the recombination is inhibited, and the dramatic improvement of dye degradation efficiency was reported in the coupled system [25, 26]. As an alternative solution to eliminate dye, the degradation of dyes by the irradiation of visible light, which is commonly referred to photosensitization, is proposed [27, 28]. In this process, the dye which is adsorbed on photocatalyst is excited by visible light and then the charge from the excited dye is injected into the conduction band of the TiO2. The oxidized form of the dye can then undergo further degradation as follows: dye+hλ→dye* dye*+TiO2→dye*++TiO2(e) TiO2(e)+O2→O2・dye*+→products dye*++ O2・-→products
(1-1) (1-2) (1-3) (1-4) (1-5)
This excellent character is advantageous to the decomposition of dyes by solar light alternative to UV irradiation [29]. However, the higher concentration of dyes often inhibited the degradation efficiency by solar light [30, 31].
1.4. Advanced biological treatment Ozonation and AOPs are useful in terms of the rapid degradation of dyes. However, the high operation cost of these processes is major drawbacks. In addition, the combination with subsequent treatment such as biological treatment is needed because the complete mineralization is not observed in only ozonation or AOPs [17, 27, 32-34]. Biological methods are generally considered environmentally friendly, because they can 6
lead to complete mineralization of organic pollutants at low cost. Though xenobiotic compounds such as azo dyes can be expected to be recalcitrant to biodegradation, it is now known that several microorganisms, including fungi, bacteria, yeasts and algae can decolorize and even completely mineralize many azo dyes under certain environmental conditions [35, 36]. The utilization of these microorganisms may lead to advanced biological treatment methods.
1.4.1. Fungal degradation of dyes based on peroxidases White-rot-fungi and other fungi can degrade various persistent organic pollutants through oxidative mechanism. The isolated strain, Thanatephorous cucumeris Dec1 could degrade various dyes, and produced a novel family of heme peroxidase, DyP, characterized by broad dye decolorization activity [37]. As the mass production of DyP was performed by gene recombinant technology, the research interest was focused not only on the potential use in bioremediation in wastewater contaminated with synthetic dyes, but also on the structural analysis of the novel enzyme [38]. The wood-rotting fongus, Phanerochaete chrysosporium in particular has been subject of intensive research related to the degradation of a wide range of recalcitrant xenobiotic compounds, including azo dyes [4]. Color removal involved lignin peroxidase (LiP) and Mn-dependent peroxidase or laccase [39]. However, the key enzyme for dye degradation, LiP, was induced under either carbon or nitrogen limitation, and the presence of carbon and nitrogen compounds found commonly in industrial effluents suppressed the production of LiP. Dye degradation by LiP consumed a considerable amount of hydrogen peroxide and veratryl alchol as reagents. Therefore, the supplementation of these reagents will be needed [40]. In continuous use of fungi,
7
mycelia aggregates clogged reactor after a short time [41]. Thus, technical advancement is needed for its applicaton to actual wastewater treatment.
1.4.2. Anaerobic decolorization by methanogenic granular sludge Azo dyes are generally persistent under aerobic condition [7, 42]. However, under anaerobic conditions, they undergo relatively easy reductive fission, yielding colorless aromatic amines. Upward-flow anaerobic sludge bed reactor (UASB), which has often been used for food industry, exerted the decolorization of azo dyes with a high degree of efficiency and stability [43]. The granules, which are the cell aggregates of anaerobes such as methanogenic bacteria and acidogenic bacteria in the biofilms containing filamentous polysaccharide slime, were formed in UASB. Methanogen degraded organic pollutants into methane in cooperation with acidogen and the retention of high densities of methanogen in UASB was suitable for the high-rate treatment of organic materials. In addition, the biofilms strengthened the resistance to xenobiotics compounds such as azo dyes [43, 44]. Twenty types of azo dyes were completely decolorized by the granular sludge from UASB [45]. Anaerobic dye decolorization appears to be nonspecific and requires co-substrates like glucose, propionate, acetate, ethanol, and so on. The decolorization is a fortuitous process, where dye might act as an acceptor of electrons supplied by the oxidation of co-substrates [36]. In this process, the electron transfer to azo dyes controlls the decolorization rate. The addition of redox medeators such as hydroquinone and anthraquinone dramatically enhanced the decolorozation efficiencies of UASB because of the electron shuttling effect [44]. The operation condition of UASB such as dye loading rates, pH, temperature and the co-substrate selection has been studied not only in laboratory scale but also in pilot
8
scale. As textile effluents often have high organic load like starch, greases, acetic acid and so on, the simultaneous removal of dye color and COD was advantage of UASB [46]. However, aromatic amines produced from azo dyes mostly cannot be degraded under anaerobic condition. It was reported that a few examples of aromatic amines could be degrade completely under anaerobic condition after 100 days [43]. Therefore, anaerobic treatment was required to combine with conventional aerobic treatment, such as sequential UASB-aerobic treatment which is feasible to dye effluent because aromatic amines are easily degraded under aerobic condition. However, the detection of aromatic amines is difficult by HPLC because of auto-oxidation of aromatic amines [47, 48]. The decrease or disappearance of aromatic amines was sometimes detected by diazotization-based method [49, 50], but it was not applied to colored samples [51]. It is possible that some aromatic amines are converted into highly colored compounds such as dimerization through auto-oxidation [48].
1.4.3. Aerobic bacterial decolorization and degradation Some bacteria decolorize azo dyes by azoreductases which catalyze the reduction of azo dyes to aromatic amines. NADH is a carrier of electron generated from the oxidation of carbon source and chemical redox compounds like different quinones sometimes mediate the reduction of azo dyes as shown in Fig. 1-2. These bacteria are often facultative anaerobes and the decolorizations are observed under semi-aerobic (static) culture. The strain Bacillus OY 1-2 were applied to the textile effluent in Wakayama city, and the effluent achieved the necessary municipal color consent limits along with a cut of sludge amounts and without reagent costs such as coagulants [52].
9
Azo dye was completely decolorized and mineralized by the anaerobic-aerobic sequential culture of the bacterial community [53]. The biofilm drum reactor mineralized the Orange II as the sole carbon source under aerobic condition, where Orange II was firstly subjected to the reductive cleavage of azo bond and then aromatic amined were mineralized. The isolated strain, Sphingomonas sp. 1CX degraded azo bond of Orange II and mineralized 1-amino-2-naphthol produced from Orange II [54]. Pseudomonas aeruginosa and Shewanera decolorationis could decolorize azo dyes under reductive mechanism and mineralize aromatic amines which had naphthol structure [55-57]. Many of constituent aromatic amines from water-soluble azo dyes have been sulfonated, and polar structures make them difficult for bacteria to uptake the aromatic amines [58].
For example, biodegradation of simple sulfonated aromatic
amine, aminobenzene sulfonic acid (ABS) is limited to a few reports [59-62]. The aerobic biodegradation of ABS had to be stimulated via bioaugmentation of an ABS-enrichment culture [63]. The addition of specifically adapted bacteria was suggested for the total mineralization of sulfonated azo dyes [64]. For the efficient decolorization process, the development of amine-metabolizing bacterial consortium was desired [64].
1.5. The use of zero-valence metals for environmental contaminants The usage of zero-valence metal (ZVM) for the remediation of environmental pollutants has recently been considered to be attractive approach because of the ease to handle and relatively low-cost [65-67].
ZVMs such as iron, tin, and zinc are
moderately strong reducing agents that are capable of reducing many common environmental contaminants in water, for example nitrate [68], nitrobenzene [69, 70],
10
pesticides [71], and halogenated organics [67, 72]. Table 1-3 shows standard electrode potentials of some metals [73]. These metals which have negative potentials are oxidized to divalent ion accompanied with the reduction of hydrogen ion or water, and produce metal ions and hydrogen gas. M0→M2++2e2H++2e-→H2 2H2O+2e-→H2+2OH-
(1-6) (1-7) (1-8)
If the contaminants are present in water/metal system, the contaminants will be reduced, competing with the reduction of hydrogen ion and water. The choice of metal is important in terms of the less influence of metal/water reactivity to metal/pollutants reactivity. Magnesium has highly negative reducing potential, and thus high efficiency of the reduction of pollutants is expected. However, magnesium/water reaction is considerably intense and predominant to the reduction of pollutants [74]. Iron, zinc and tin have mild reducing potential metal, and are better choice since the metal/water reaction is lower and less influential to the metal/pollutants reactions. Zinc has highest reducing power among three metals and was more effective for dechlorination of chlorinated organic pollutants than other metals [72, 74]. However, the release of Zn (II) species can be serious concern in water contamination [67]. Iron has been used most popularly for the reduction of pollutants, since it is relatively inexpensive and environmentally acceptable. Table 1-4 shows the environmental contaminants which have been treated by iron [65]. Recent investigation on dechlorination of chlorinated organic pollutants by zero-valence iron has shown promising potential for applying the technology to the in situ remediation of contaminated groundwater. Numerous feasibility studies, pilot tests, and field scale demonstrating projects have been carried out using granular iron [75-78]. 11
1.5.1. The reduction of azo dyes by zero valence iron Zero-valence iron has been studied to decolorize textile wastewaters, especially in China, where the production of synthetic dyes is the largest in the world [79-84]. Iron can decolorize azo dyes through reductive cleavage of azo bonds. As hazardous aromatic amines remained in treated water, post treatment like activated sludge was required. The iron pretreatment of azo dyes is considered to be viable option for the subsequent aerobic biological treatment [85]. The decolorization kinetics by iron depended mainly on surface area of iron, solution pH and initial dye concentrations [79, 81, 86]. The decolorization of azo dye occurred on the surface of iron [87], and thus the reaction rate is proportional to the surface area [79]. The azo dye molecules received electrons from the iron and combined with H+ from an acid to form the transitional product. This product gained electrons and combined with H+ again, forming the terminal product. Its reaction mechanism is shown in Fig. 1-3. The efficiency of decolorization increased as the solution pH decreased. This is not only because hydrogen ion took part as one reactive specie in the decolorization reaction, but also because the reduction potential of the reaction Fe0-2e=Fe2+ increased with decreasing in the solution pH [81, 88]. The decolorization reaction by zero-valence iron appears to be the simple oxidation-reduction reaction, but the exact mechanism is not fully understood. The decolorization reaction includes the heterogeneous reactions such as diffusion, adsorption of dyes to the surface of iron, and the surface reaction of dyes. In addition, the oxide film formation of the iron surface affects the decolorization efficiency. These phenomena make it difficult to analyze the kinetic data and give us interesting research topics.
12
Though laboratory scale experiments have been extensively performed, zero-valence iron processes have some drawbacks in the application to actual wastewater treatment. Firstly, the surface of iron was easily covered with oxide film in the presence of dissolved oxygen to inhibit the reactivity. Secondly, the huge amounts of colored precipitates were produced by rusting.
1.5.2. The combined use of zero-valence iron and biological treatment The feasibility of the combined use of zero-valence iron and anaerobic microorganisms has been demonstrated. Some examples are that microbial mediated dechlorination of chlorinated organic pollutants with the combinated use of zero-valence iron (ZVI) and anaerobic microorganisms [89-91], the conversion of nitrobenezene into aniline by anaerobic sludge accelerated by adding zero-valence iron [92], and a zero-valence iron packed upward flow anaerobic sludge blanket reactor, ZVI-UASB to enhance azo dye wastewater treatment [93, 94]. However, the treatment of aromatic amines after the decolorization is required. There are only a few published papers regarding aerobic biodegradation of aromatic amines which were produced after iron-pretreatment of azo dyes [85, 95, 96]. Aniline and sufonic acid (p-aminobenzene sulfonic acid) produced from iron treatment of Orange G and Orange I respectively were degraded in acclimated activated sludge cultures. However, the iron-pretreatment has to be conducted under strict anaerobic condition to avoid rusting of iron. The strictly anaerobic operation will give difficulty in actual system and it is unknown whether the naphthol structured aromatic amines as second products were biodegraded or not because these aromatic amines are known to undergo auto-oxidation reaction and an analytical standard is not available [47].
13
1.6. Purpose of this study The treatment of azo dyes by zero-valence iron has been extensively studied [83, 84]. However, in situ application of zero-valence iron is limited by its rusting to produce huge colored precipitates. The use of zero-valence tin may be able to avoid the problem. In contrast to iron, there is almost no report with respect to the decolorization of azo dyes by tin. Its potential application in the decolorization of wastewaters from the use of dyes is not known. In this study, the basic studies regarding the decolorization of azo dyes by zero-valence tin and tin immobilized material, PIP (Powder impact plating) tin ball, were conducted and the subsequent biological treatment in terms of the mineralization of produced aromatic amines was conducted. Also the decolorization of an azo dye, Orange II under aerobic condition was investigated by the combined use of zero-valence tin and Orange II decolorizing bacterial strain, Pseudomonas aeruginosa m2. Zero-valence tin dissolves as bivalent ions and subsequently is oxidized to tetravalent ions in strong acid or alkaline solutions [97, 98]. Citrate ion is known to promote the dissolution of tin due to the formation of the stable aqueous chelate between Sn2+ ion and citrate [99]. In similar way, carbonate accelerates the dissolution of zero-valence iron [69, 100-102]. Zero-valence tin /citric acid system is also expected to be the potential application to the reduction of pollutants. Firstly, the purpose of this study is to elucidate the feasibility of the decoloriztion of azo dyes by zero-valence tin especially in citric acid solution. Secondly, the combination of zero-valence tin and biological treatment was planned for the following purposes: 1) biodegradation of aromatic amines produced after the tin treatment, and 2) the achievement of decolorization at neutral pH under aerobic condition.
14
In chapter 2, the title is “Decolorization of azo dyes by PIP tin ball in citric acid solution”. Tin immobilized material, PIP tin ball, was used for the decolorization of azo dyes as a novel heterogeneous process in the presence of citric acid and in the absence of citric acid. The decolorization rates in PIP tin ball/ citric acid system were estimated under aerobic condition and anaerobic condition and the repeated decolorizations of Orange II were performed. XPS analysis of repeatedly used PIP tin ball and unused PIP tin ball were performed to clarify the reason for the inactivation of PIP tin ball. In chapter 3, the title is “Comparative study on the decolorizations of Orange II by zero-valence tin in citric acid and in hydrochloric acid solutions”. Chapter 2 implied that the oxidation of zero-valence tin in the structure PIP tin ball was considered to be a role for the reaction in citric acid. The metallic tin was used for the decolorization of Orange II to investigate the interaction between zero-valence tin and citric acid. The decoloization kinetics of Orange II in citric acid were compared with those in hydrochloric acid. The enhancement effect of citric acid in the decolorization reaction of Orange II by zero-valence tin was discussed. In chapter 4, the title is “Biodegradation of two aromatic amines produced from the decolorization of Orange II by zero-valence tin”. To degrade two aromatic amines, p-aminobenzene sulfonic acid (ABS) and 1-amino-2-naphthol (1A2N), which were produced from tin treatment of Orange II, the enrichment cultures were performed. Characterization of the isolated strains with respect to the degradation activity against aromatic amines was carried out. The sequential treatments of the decolorization of Orange II by zero-valence tin and bacterial communities of isolated strains were performed. In chapter 5, the title is “Decolorization of Orange II by combined use of tin and
15
Pseudomonas aeruginosa m2 in the presence of tin”. P. aeruginosa m2 strain, which was isolated as a degrader of 1A2N, could decolorize Orange II in the presence of citrate under static condition. The decolorization of Orange II by combined use of zero-valence tin and P. aeruginosa m2 were performed at neutral pH and under aerobic condition, where the decolorizations did not occurred in the individual use. Chapter 6 is the conclusions which summarized the results in this study, and compared with other treatments and the future aspect were described.
16
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24
Table 1-1. Estimated degree of fixation for different dye-fibre combinations and loss to effluent [3]
Dye application classes Acid Basic Direct Disperse Metal-complex Reactive Sulfer Vat
Degree of fixation (%) 89-95 95-100 70-95 90-100 90-98 50-90 60-90 80-95
Fibers Polyamido Acryl Cellulose Polyester Wool Cellulose Cellulose Cellulose
25
Loss to effluent (%) 5-20 0-5 5-30 0-10 2-10 10-50 10-40 5-20
Table 1-2. Advantages and disadvantages of representative promising methods of dye removal. Physicochemical process Methods Ozone
Advantage Applied in gaseous state: No alternation of volume Effective decolorization besides mineralization
Disadvantage Short half-life (20 min)
TiO2/UV
The high stability of oxide titanium
Fast recombination of electron and hole
Zero-valence iron
The fast conversion of azo dyes to biodegradable aromatic amines
Rusting
Fenton
Sludge production
Advanced biological processes Methods Fungi
Advantage Aerobic degradation
Disadvantage Long term for enzyme production
UASB
Applied in large scale treatment
Sensitivity of methanogenic bacteria
(+aerobic treatment) Facultative anaerobic bacteria (+aerobic treatment)
The high degree of decolorization and degradation of azo dyes in a single reactor
(Difficulty in the detection of aromatic amines) Almost no information regarding application to actual system (Difficulty in the construction of effective bacterial communities)
The combination of physicochemical process and biological process Methods Ozone or AOPs*+ Zero-valence iron+
Advantage The high performance to actual system The effective decolorization process is constructed besides the degradation of produced aromatic amines
*
AOPs: Advanced oxidation processes
26
Disadvantage The possibility of no biodegradable products after AOPs* No information regarding application to actual system
Table 1-3. Standard electrode potentials of bivalent metal/metal ion half-reactions [73]. M2+/M Ba +2 e-→Ba Sr2++2 e-→Sr Ca2++2 e-→Ca Mg2++2e-→Mg Mn2++2 e-→Mn Zn2++2 e-→Zn Fe2++2 e-→Fe Cd2++2 e-→Cd Co2++2 e-→Co Sn2++2 e-→Sn Pb2++2 e-→Pb ( 2H++2 e-→H2) Cu2++2 e-→Cu Hg2++2 e-→Hg Pt2++2 e-→Pt 2+
Standard electrode potential (V vs SHE) -2.925 -2.89 -2.84 -2.356 -1.18 -0.763 -0.44 -0.403 -2.77 -0.138 -0.126 (0) 0.337 0.85 1.188
27
Table 1-4. Contaminants treated by zero-valent iron [65]. Organic compounds Methane
Ethanes
Ethenes
Propanes Aromatics
Others
Inorganic compounds
Tetrachloromethane Trace metals Trichloromethane Dichloromethane Hexachloroethane 1,1,1-trichloroethane 1,1,2-trichloroethane 1,1-dichloroethane 1,2-dibromoethane Tetrachloroethene Trichloroethane cis-1,2-dichloroethene trans-1,2-dichloroethene Anions Vinyl chloride 1,2,3-trichloroprpane 1,2-dichloropropane Benzene Toluene Etylbenzene Hexaachlorobutradiene Freon 113 n-nitrosodimethylamine
28
Arsein Cadmium Chromium Cobalt Copper Lead Manganese Nickel Selenium Uranium Zinc Nitrate Phosphate Sulphate
UV/H2O2
: H2O2 + hλ → 2HO・
Ozone/UV
: O3 + hλ → O2 + O O + H2O → 2HO・
Ozone/H2O2 :
H2O2 + H2 ⇔ HO2- + H+ HO2- + O3 → HO・+ O2- +O2
Fenton reagent : Fe2+ + H2O2 → Fe3+ + OH- + HO・ Fe 3+ + H2O2 → Fe2+ + HO2・ + H+ O2 Conduction band
electron
O2-
TiO2/UV :
Valence band
hole
H2O OHHO・
Fig. 1-1. The formation mechanism of hydroxyl radical (HO・) by various AOPs.
29
Colorless solution containing amines
Colored solution containing azo dye X
X
Redox MediatoroX
N
Redox Mediatorred
H2N
Azoreductase
N NAD+
NADH
NH2 Oxidation Products
Carbon Complexes
Dehydrogenase (one electron oxidation)
X Cell
X
Fig. 1-2. Proposed mechanism for reduction of azo dyes by whole bacterial cells.
30
HO
HO Fe(0)
N
HO 3S
N +
Fe HO 3S
2H
N
N
H
H
HO Fe(0) +
HO 3S
NH2
+
H2N
2H
Fig. 1-3. Degradation mechanism of Orange II in Fe(0) –H2O system
31
2+
Chapter 2 Decolorization of azo dyes by PIP tin balls in citric acid solution
2.1. Introduction The treatment of dye wastewater has attracted attention because of local landscape impacts and also because of potential health effects to living cells. Azo dyes account for 50% of total dyes produced worldwide and have potential risk because of toxic aromatic amines produced as reductive metabolic intermediates. Although the physicochemical treatments such as coagulation, flocculation is a viable option for the elimination of azo dyes, these produce huge amounts of sludge which pose handling and disposal problems. Microbial treatment of azo dyes has been reported in a variety of reaction systems [1-3]. Azo dyes are degraded to aromatic amines under anaerobic conditions, but further degradation of these compounds under anaerobic conditions does not occur. As aromatic amines are decomposed under the aerobic conditions, the use of anaerobic-aerobic sequential batch systems has been proposed [2, 4, 5]. However, the decolorization efficiency of azo dyes remains unsatisfactory mainly because of poor decolorization reaction rates. The utilization of heterogeneous zero-valence iron powders for pretreatment of azo dyes prior to aerobic biological processes is an alternative approach. [6-8]. As novel heterogeneous process, I proposed the decolorization of azo dyes using immobilized tin components produced by powder impact plating (PIP tin balls). In contrast to iron, there is almost no report with respect to the decolorization of azo dyes using tin. Its potential application in the decolorization of wastewaters from dye use was not known. PIP tin ball is commercially available, and has been utilized to purify the atmosphere indoor. The reaction mechanism is not yet entirely clear. In heterogeneous 32
process, the separation of solid materials after the treatment is important factor with respect to wastewater treatment. PIP tin ball has comparatively large volume, which is easy to recover. Additionally, vigorous agitation is not needed, different from the type of suspended particle media because of the large effective volumetric capacity, which is costly effective point. If the usage of PIP tin balls for decontamination of dye waste is possible, novel effective treatment could be proposed. It was found that PIP tin balls decolorized four azo dyes which are widely used in textile industries under static condition in which no stirring is needed and at room temperature. I mainly investigated the decolorization characteristics of Orange II as a model azo dye. The decolorization was enhanced in the presence of citric acid. I proposed the decolorization reaction mechanism of azo dyes by PIP tin balls in the presence of citric acid.
2.2. Methods 2.2.1. Materials The chemical structures of Orange II, Acid yellow 23 and Acid red 27 (Tokyo Kasei Kogyo Co., Ltd., Tokyo), and Reactive black 5 (Nippon Kayaku Co., Ltd., Tokyo) are shown in Fig. 2-1. These reagents were used without further purification. The reagents, p-aminobenzene sulfonic acid (ABS) and citric acid (Kanto Chemical Co., Ltd., Tokyo) and 1-amino-2-naphthol (1A2N) (Wako Pure Chemical Industries Co., Ltd., Tokyo) were of guaranteed reagent grade. The basic structure of PIP tin balls (4 mm diameter; BET surface area 0.0035 m2/g; average density 3.5876 g/cm3; Fuji Kihan Co., Ltd., Nagoya) is shown in Fig. 2-2. The PIP tin balls consist of an inner carrier of Al2O3, a tin layer around the carrier and an outer surface of tin oxide (SnOx, where x is the stoichiometric ratio of oxygen).
33
2.2.2. Decolorization of azo dyes under static conditions Decolorization reactions were carried out by addition of 2.2 ml of different concentrations of azo dyes to Pyrex test tubes (1 cm diameter, 10 cm long) containing 8 g of PIP tin balls under static conditions at room temperature and atmospheric pressure. The total volume of 8 g PIP tin balls was about 2.2 cm3. The decolorization was performed at pH 3 in the presence of citrate buffer or hydrochloric acids. The maximum absorbance of each dye (OrangeⅡ:485nm, Acid red 27:504nm, Acid yellow 23:426nm, Reactive black 5:595nm) was measured using a UV-Visible spectrophotometer (UV-2200 Shimadzu, Kyoto), and absorbance was proportional to concentration over the tested range. The effect of varying citric acid concentrations (0-15 mM) on the decolorization ratio was also investigated during 7 min reaction time under static conditions. In this case, pH was not controlled. Decolorization ratios were estimated as follows. Decolorization ratio (%)= (C – CPIP)/C×100
(2-1)
where C is the concentration of OrangeⅡin the case without PIP tin balls after 7 min and CPIP is the concentration of OrangeⅡin the case with PIP tin balls after 7 min. The initial reaction rates at various initial concentrations of dyes were determined. All experiments were carried out two times, and the average values were used.
2.2.3. HPLC analysis and product identification Degradation products of Orange II after static treatment with PIP tin balls in 5 mM citric acid were analyzed by high performance liquid chromatography (HPLC) using filtered reaction solutions through membrane filter (pore size 0.2 μm; Advantec, Tokyo). 34
The HPLC was a Jasco PU-980 delivery system equipped with a UV-VIS detector (875UV Jasco, Tokyo), a column oven (TU-100 Jasco) and a reverse phase C18 column (AG 120, 0.46 cm×15 cm; Shiseido, Tokyo). Elution conditions were as follows. The first step was under isocratic conditions with 0.1 M phosphate buffer (pH 7) for 5 min, the second step used a 10 min linear gradient of water : methanol (1:1 v/v) mixture from 0％ to 100％ and the final step employed an isocratic gradient of water : methanol (1:1 v/v) mixture at a flow rate of 1.0 mL/min for 10 min. Eluted materials were detected by absorbance at 254 nm. p-Aminobenzene sulfonic acid (ABS) was identified by HPLC-mass spectrometry (LC-MS;LC:2690, ESI-MS; ZQ, Waters, Melford). HPLC conditions were the same as described above. 1-Amino-2-naphthol (1A2N) was identified by comparing the retention time of the material with that of an authentic sample, and by co-elution of the test compounds with standard samples. Quantification of products was performed using calibration curves prepared from HPLC of authentic samples.
2.2.4. Decolorization of OrangeⅡunder stirred condition Reaction solution containing 0.1 mM Orange II and 5 mM citric acid was deoxygenated by bubbling nitrogen gas for 20 min. PIP tin balls (20 g) were placed in an EVA (ethylene vinyl acetate copolymer) mesh bag which is non-reactive to PIP tin balls and seldom adsorbs dye, and then 300 ml of the deoxygenized reaction solution was added to the 500 ml flask which equipped EVA mesh bag entrapping PIP tin balls. Nitrogen gas was bubbled into the solution at room temperature and the mixture was stirred using a magnetic stirrer throughout the reaction. Similar decolorization reactions with continuous stirring were carried out under bubbled air conditions and under 35
ambient conditions. The gas flow rate was 10 ml/min. The potion of 1ml was sampled out
periodically.
The
concentration
of
Orange Ⅱ was
measured
by
the
spectrophotometer.
2.2.5. Repeated decolorization of OrangeⅡunder anaerobic condition PIP tin balls (7 g) were placed in an EVA mesh bag suspended in a 100 ml round-bottom flask. The flask was sealed with a silicone stopper equipped with a three-way valve, a two-way valve and a sampling tube. Two ports of the three-way valve were connected to a vacuum pump (pump port) and a balloon filled with nitrogen gas (balloon port), respectively. To maintain nitrogen atmosphere in a flask one port of the two-way valve was connected to a nitrogen cylinder which provided a continuous nitrogen supply. Complete replacement of air in the flask with nitrogen gas was achieved as follows. The balloon port of the three-way valve, the two-way valve and the sampling tube were closed, and the flask was evacuated. The pump port was then closed, and the balloon port was opened to supply nitrogen to the flask. This was repeated three times and then the balloon was removed, and nitrogen gas was supplied continuously. Reaction solution (100 ml of 0.1 mM OrangeⅡ in 5 mM citric acid) was deoxygenated by continuous bubbling with nitrogen gas for 20 min and then added to the flask with a syringe from the balloon port. After 50 min, 60 ml of reacted solution was removed from the flask, 60 ml of fresh deoxygenated dye solution was added, and decolorization was continued. This procedure was repeated five times. The experiment was also conducted under ambient conditions. In this case, the deoxygenation procedure of reaction solution and continuous nitrogen supply was eliminated, and two-way valve was opened to expose reaction solution to air during reaction. At each repeated cycle, 36
initial concentration and the concentration of OrangeⅡafter 50 min was measured by the spectrophotometer. The decolorization ratio at 50 min was compared at each cycle. The decolorization ratio was estimated as follows. Decolorization ratio (％)= (Co-C50min)/Co×100
(2-2)
, where Co is initial concentration and C50min is the concentration of OrangeⅡafter 50 min.
2.2.6. FT-IR measurement Fourier transform infrared spectrometry (FT-IR) was used to analyze the white precipitate which appeared on inactivated PIP tin balls during the experiment under ambient conditions. The surface of the balls was shaven by a blade to remove the precipitate. A disk including the precipitate was prepared by pressing the shavings with KBr powder and used for FT-IR (660-plus; Jasco). Crystalline citric acid was used as reference.
2.2.7. XPS analysis X ray photoelectron spectroscopy (XPS) spectra of fresh and inactivated PIP tin balls were obtained at Shimadzu Co., Ltd. Laboratory. The non-mono-chromatized Al-Kα radiation source was applied to two PIP tin balls at 15 kV and 10 mA. The measurement was performed with a take-off angle close to 90o.
2.3. RESULTS 2.3.1. Decolorization of azo dyes by PIP tin balls under static conditions Table 2-1 shows the decolorization ratios for four azo dyes (all at 0.3 mM) after 7 min 37
reaction time under static conditions using PIP tin balls in citrate buffer (5 mM, pH 3) or in hydrochloric acid at pH 3. Decolorization ratios in the presence of citrate were greater than those in hydrochloric acid, except for Reactive black 5. Figure 2-3 shows the decolorization ratios of Orange II by PIP tin balls with a reaction time of 7 min under static conditions and solution pH at various citric acid concentrations. The decolorization reaction was not occurred without PIP tin balls. The decolorization ratio increased to 70% in a concentration-dependent manner with increasing citric acid concentration up to 10 mM, but no increase in the decolorization ratio was observed above this concentration.
2.3.2. Products after decolorization of Orange II HPLC chromatogram of the products of Orange II degradation after 10 min under static conditions in 5 mM citric acid with PIP tin balls showed three major peaks. Orange II was detected at a retention time of 30.1 min. Two new peaks appeared with retention times of 2.1 and 20.6 min as shown in Fig. 2-4. Figure 2-5 shows the negative ion MS spectrum of the component with a retention time of 2.1 min. The ion peak (M-H-) at 190.7 is indicative of citric acid (molecular weight 191). An ion peak at (M-H-) 171.7 was confirmed to be p-aminobenzene sulfonic acid (ABS) (molecular weight 173) by correlating the HPLC retention time with that of an authentic sample. The component with the retention time of 20.6 min was identified as 1-amino-2-naphthol (1A2N) by comparing the retention time of the sample with that of an authentic sample and by co-elution by HPLC. HPLC and LC-MS results indicated that Orange II was reductively cleaved by PIP tin balls to the two hydrogenated products, ABS and 1A2N as shown in Fig. 2-6.
38
Time-course concentrations of Orange II and the two degradation products during reaction with PIP tin balls are shown in Fig. 2-7. Rapid analysis after sampling of reaction solution and the acidic solution could avoid the effect of auto-oxidation of 1A2N on the quantification. Orange II was not detected after 20 min, at which time stoichiometrically equivalent amounts of ABS and 1A2N accumulated, and their concentrations showed no further change after 30 min. This shows that the degradation of Orange II by PIP tin balls was irreversible, and that further decomposition of the two products did not occur.
2.3.3. Effect of oxygen on the decolorization of Orange II The time course of Orange II degradation by PIP tin balls under no gas supply (ambient), air bubbling and nitrogen bubbling conditions was investigated. A plot of ln (Co/C) versus time showed a linear relationship (Fig. 2-8), indicating the kinetics fit pseudo first-order model. The reaction rate constant obtained under nitrogen bubbling conditions was 5.1×10-2 (min-1), which was greater than with under air bubbling conditions (2.4×10-2 min-1) or no gas supply conditions (2.9×10-2 min-1).
2.3.4. Repeated decolorization of OrangeⅡ Decolorization in repeated 50 min reactions was performed under anaerobic conditions established by continuous bubbling with nitrogen. A control experiment was conducted under ambient conditions. The initial concentration of Orange II and the decolorization ratios obtained during five reaction cycles are shown in Table 2-2. Almost complete decolorization was obtained by 50 min reaction in both cases. A decrease in the
39
decolorization ratios under ambient conditions was evident, but only a slight decrease in decolorization ratios was observed under anaerobic conditions. A white precipitate progressively covered the surface of the tin balls during repeated decolorization under ambient conditions. Once the white precipitate was removed, the ball regained decolorization activity (date not shown). This indicates that the precipitate on the surface of the balls was related with the inactivation of the balls. Fresh (unused) PIP tin balls and the PIP tin balls which were inactivated by repeated use under ambient condition were analyzed by XPS. Sn3d spectra of unused PIP tin balls and an inactivated ball are shown in Figure 2-9. The small bands at 485 eV (b in the figure) and 493 eV (a in the figure) in the spectrum of unused PIP tin balls were assigned to metallic tin and were not found in spectrum of inactivated PIP tin balls. According to Willemen et al., the center positions of the Sn 3d5/2 spectrum peaks in Sn(Ⅱ) and Sn(Ⅳ) oxide compounds are 486.3±0.6 eV (SnO) and 486.9±0.6 eV (SnO2), respectively [9]. This means that it is not possible to distinguish Sn(Ⅱ) and Sn(Ⅳ) from the binding energy data. The center position of the Sn 3d5/2 of inactivated one was slightly higher compared with that of fresh one. This indicated that the tin of higher oxidation state is occupied in inactivated PIP tin balls. Atomic concentrations of C, O, and Sn were investigated in fresh and inactivated PIP tin balls from XPS. Atomic ratios to total Sn are shown in Table 2-3. The increase in the atomic ratios of C/Sn and O/Sn was observed in inactivated PIP tin balls. This indicates that the precipitates covered the surface of inactivated PIP tin balls are Sn(Ⅳ) species such as SnO2 and/or Sn(OH)4 along with incorporating carbon species. FT-IR spectra of the precipitate and citric acid monohydrate were obtained and compared with those reported by Tselesh [10] for Sn(Ⅱ)-citrate chelate, SnO2・nH2O 40
and Sn(OH)2 as shown in Table 2-4. Bands at 1390 cm-1 which were attributed to carboxylate COO- were found in the spectra of the precipitate, citric acid monohydrate and Sn(II)-citrate. Bands at 1723 cm-1 and 1228 cm-1 in the spectra of citric acid monohydrate which were attributed to stretching of the C=O double bond and of the C-O single bond in carboxylic acid, respectively [11] were not found in spectra of the precipitate and Sn(II)-citrate. This indicates that the carboxylic group in the precipitate or Sn(II)-citrate may be chelated. As Sn(II) forms a stable complex with citrate [10, 12], it is assumed that dissolved Sn (Ⅱ) from PIP tin balls also forms a complex with citrate. A set of absorption at 2921, 2851 cm-1, assignable to ν (CH) vibration of CH2 groups of citrate ions [10] was exhibited in precipitates, while the absorptions were not found in citric acid monohydrate. The absorption bands at 1136 cm-1 in spectra of the precipitate and Sn(II)-citrate, and 1158 cm-1 in the spectrum of citric acid monohydrate corresponded to bending vibrations of δ-(OH) of C-OH groups. The peak at 1245 cm-1 of the spectrum of the precipitate attributed to terminal Sn-OH stretching was also found in Sn(II)-citrate, SnO2・nH2O and Sn(OH)2. The bending vibration of O-Sn at 659 and 549 cm-1 in spectrum of the precipitate was found for Sn(II)-citrate, SnO2・nH2O and Sn(OH)2. These results indicated the possibility that tin-citrate complex was included in precipitates. Additionally, it is assumed that the tin-citrate complex became poorly soluble by combining tin oxide and/or tin hydroxide species.
2.4. Discussion PIP tin ball decolorized azo dyes and the decolorization reactions ware enhanced in the presence of citric acid. Decolorization products of Orange II were identified as p-aminobenzene sulphonic acid and 1-amino-2-naphthol which were produced by 41
reductive cleavage of azo group. Considering subsequent biological treatment, complete decomposition of OrangeⅡ is needed. It is reported that OrangeⅡ inhibits microbial metabolism [13]. Under nitrogen bubbling conditions, almost complete decolorization was
obtained
by
50
min
reaction
under
stirring
condition
and
the
surface-area-normalized first order rate constants (kSA) was estimated to be 0.215 (l m-2 min-1), which is almost the same as the reported kSA=0.21 (l m-2 min-1), for decolorization of 0.3 mM OrangeⅡwith zero-valence iron [7]. XPS analysis showed that metallic tin Sn(0) on the surface of the fresh PIP tin ball was not found on inactivated one. This indicates that metallic tin, Sn(0) as a reducing agent is converted to Sn(II) by producing electron which plays a role of the decolorization of azo dyes. The decolorization of four azo dyes was investigated in the presence of citrate and in the absence of citrate. It is reported that the order of the decolorization rate in microbial reduction was Acid red 27, OrangeⅡ, and Acid yellow 23, in which the microbial azo reduction occurs sequentially as a function of the substrate reduction potential [14,15]. To investigate the correlation in PIP tin balls, three azo dyes were selected. The correlation between the reduction potentials of these three dyes and these decolorization ratios were not found in PIP tin balls (Table 2-1). I also used Reactive black 5 which belongs to a class of compounds known as reactive azo, which are abundantly used in textile industries for dyeing [16]. The decolorization ratio of OrangeⅡwas the highest among these dyes (Table 2-1). In Acid red 27 and Acid yellow 23, dramatic enhancement of ratios was observed in the presence of citrate. OrangeⅡ and Reactive black
5
showed
obvious
decolorization
even
by
hydrochloric
acid.
Langmuir-Hinshelwood model was applied in order to study the adsorption of dye on 42
the surface of PIP tin balls to calculate the equilibrium constants. The reaction rate of surface processes could be described by the Langmuir-Hinshelwood equation where rate is proportional to the surface coverage of reactive active sites. The equilibrium constants were determined by fitting the experimental date under static condition to the following Langmuir-Hinshelwood equation: 1/vo=1/kapp+1/(kappKm[Co])
(2-3)
,where vo [mM/min], kapp [mM/min], Km [mM-1], and Co [mM] are initial reaction rate, apparent reaction rate constant, adsorption equilibrium constant, and initial dye concentration in the decolorization reaction, respectively. A linear expression can be obtained by plotting the reciprocal initial rate against the reciprocal initial concentration. The kinetic parameters which were estimated from equation (2-3) are shown in Table 2-5. The adsorption equilibrium constants of Acid yellow 23 and Acid red 27 in the case with hydrochloric acid could not be determined because of these initial reaction rates were considerably small. The adsorption equilibrium constants of Reactive black 5 and OrangeⅡ were smaller value in the case including citrate than hydrochloric acid. This would be because the adsorption of azo dyes was disturbed by citrate. This disturbance would be responsible for the adsorption of citrate on the surface of PIP tin ball. The difference of adsorption equilibrium constants between two acids was considerably large in Orange II. Apparent reaction rate constant reflects the limiting rate of the reaction at maximum coverage of adsorption site. Apparent reaction rate constant of Reactive black 5 was larger in citrate than in hydrochloric acid, though the decolorization ratio was smaller with citrate as shown in Table 2-1. These results indicated that the reactive site was considerably reduced in the presence of citrate, leading to the inhibition of decolorization reaction in the case of Reactive black 5.
43
The decolorization ratios of Orange Ⅱ were investigated at various citric acid concentrations (Figure 2-3). In increasing from 2.5 mM to 5 mM, the dramatic enhancement of decolorization ratios was observed. However, the enhancement was not observed above 7.5mM, while the solution pH was gradually decreased in increasing citric acid concentrations. On the contrary, the decolorization ratios were considerably influenced by solution pH in the case of hydrochloric acid, in which higher decolorization ratio was obtained as solution pH was decreasing (date not shown). In the case of citric acid, the adsorption on citrate on PIP tin balls would be significant for decolorization reaction rather than the effect of pH. It is indicated that the oxidation process of metallic tin on the surface of PIP tin balls which is depicted in the equation (2-4) plays an important role in the decolorization reaction. Sn(0) → Sn(Ⅱ)＋2e-
(2-4)
The adsorption of citrate at active Sn(0) sites on the PIP tin balls would be followed by the oxidation of Sn(0) to Sn(Ⅱ). The oxidation reaction would be thermodynamically favored by stabilization of Sn(Ⅱ)-citrate complex [17]. It has been reported that metallic tin is oxidized to soluble Sn(Ⅱ) species in citric acid solution and a Sn(Ⅱ)-citrate complex is formed. This complex is adsorbed on the tin surface. The adsorbed Sn(Ⅱ)-citrate is oxidized and forms a more stable SnO2 film on the tin surface [12, 18]. XPS analysis indicated the increase in the ratio of Sn(Ⅳ) species in inactivated PIP tin balls. The inactivation was responsible for the precipitates formed on the surface of PIP tin ball. FT-IR analysis of the precipitates indicated that the tin-citrate complex was included in the precipitates in addition to tin oxide and/or
44
hydroxide species. The thick film covered metallic tin portion on the surface of PIP tin balls, resulting in its inactivation. As tin dissolves only slightly in acids free from air, the film formation from dissolved tin species was not so vigorously observed under anaerobic condition, leading to no inactivation [19].
45
References [1]Mielgo I., Moreria M. T., Feijoo G. and Lema J. M., A packed-bed fungal bioreactor for the continuous decolourisation of azo-dyes(OrangeⅡ), J. Biotechnol., 89: 99-106 (2001) [2]Stolz A., Basic and applied aspects in the microbial degradation of azo dyes, Appl Microbiol Biotechnol., 56: 69-80 (2001) [3]López C., Moreria M. T., Feijoo G. and Lema J. M., Dye decolorization by manganese peroxidase in an enzymatic membrane bioreactor, Biotechnol. Prog., 20: 74-81 (2004) [4]Delée W., O‟Neill C., Hawkes F. R. and Pinheiro H. M., Anaerobic treatment of textiles effluents: a Review, J. Chem. Technol. Biotechnol., 73: 323-335 (1998). [5]Van der Zee F. P. and Villaverde S., Combined anaerobic-aerobic treatment of azo dyes-A short review of bioreactor studies, Wat. Res., 39: 1423-1440 (2005) [6]Cao J., Wei L., Huang Q., Wange L. and Han S., Reducing degradation of azo dye by zero-valent iron in aqueous solution, Chemosphere, 38: 565-571 (1999) [7]Nam S. and Tratnyek P. G.. Reduction of azo dyes with zero-valent iron, Wat. Res., 34:1837-1845 (2000) [8]Mu Y., Yu H. Q., Zhang S. J. and Zheng J.-C., Kinetics of reductive degradation of OrangeⅡ in aqueous solution by zero-valent iron, J. Chem. Technol. Biotechnol., 79: 1429-1431 (2004) [9]Willemen H., Van De Vondel D. F. and Van Kelen G. P., An ESCA study of tin compounds, Inorg. Chim. Acta, 34: 175-179 (1979) [10]Tselesh A. S., Anodic behavior of tin in citrate solutions: The IR and XPS study on the composition of the passive layer, Thin Solid Films, 516: 6253-6260 (2008) [11]Silverstein R. M., Bassler G. C. and Morrill T. C., Spectrometric identification of organic compounds, Wiley, New York. (1964) [12]Almeida C. M. V. B. and Giannetti B. F., Protective film growth on tin in perchlorate and citric acid electrolytes, Mat. Chem. Phys., 69: 261-266 (2001) [13]Alquerque M. G. E., Lopes A. T., Serralheiro M. L., Novais J. M. and Pinheiro H. M., Biological sulphate reduction and redox mediator effects on azo dye decolorization in anaerobic-aerobic sequencing batch reactors, Enz. Microb. Technol., 36: 790-799 (2005) [14]Bragger J. L., Lloyd A. W., Soozandehfar S. H., Bloomfield S. F., Marriott C. and Martin G. P., Investigation into the azo reducing activity of a common colonic 46
microorganism, Int. J. Pharma., 157: 61-71 (1997) [15]Semdé R., Pierre D., Geuskens G., Devleeshouwer M. and Moës A. J., Study of some important factors involved in azo derivative reduction by Clostridium perfringens, Int. J. Pharma., 161: 45-50 (1998) [16]Işik M. and Sponza D. T., A batch kinetic study on decolorization and inhibition Reactive black 5 and Direct brown 2 in an anaerobic mixed culture, Chemosphere, 55: 119-128 (2004) [17]Survila A., Mockus Z. and Kanapeckaitė S., Kineticks of Sn and Co codeposition in citrate solutions, Electrochim. Acta, 46, 571-577 (2000) [18]Abdel Rehim S. S., Sayyah, S. M. and El Deeb, M. M., Corrosion of tin in citric acid solution and the effect of some inorganic anions, Mat. Chem. Phys., 80: 696-703 (2003) [19]Jafarian M., Gobal F., Danaee I., Biabani R. and Mahjani M. G. Electrochemical studies of the pitting corrosion of tin in citric acid solution containing Cl-, Elecrochim. Acta, 53: 4528-4536 (2008).
47
Table 2-1. The decolorization ratios for four dyes by PIP tin balls in 7 min under static conditions. The pH was adjusted to 3 using citric acid or hydrochloric acid. PIP tin ball : 8g, Initial concentration of dyes : 0.3 mM, The volume of reaction solution: 2.2 ml Decolorization ratio [%] With citrate With hydrochloric acid
Acid yellow 23
Acid red 27
Reactive black 5
Orange II
33 1
43 7
10 22
64 33
48
Table2-2. Initial concentrations of Orange II and decolorization ratios obtained in repeated 50-min reaction cycles. Repeated cycles
Anaerobic condition
Ambient condition
Initial conc. [mM]
Decolorization ratio [％]
Initial conc. [mM]
Decolorization ratio [％]
1
9.6×10－2
99
9.6×10－2
96
2
5.0×10－2
99
4.9×10－2
93
3
5.8×10－2
99
5.0×10－2
85
4
6.0×10－2
96
5.1×10－2
63
5
6.0×10－2
96
5.5×10－2
56
49
Table 2-3.
XPS quantitative results. Peak
Position [eV]
Raw Area [CPS]
Relative Sensitivity factor
Atomic ratio to total Sn
O1s Sn3d C1s
531.6 487.1 285.0
20100.0 47164.6 3922.0
0.780 4.725 0.278
2.6 1 1.40
O1s Sn3d C1s
531.8 487.4 285.1
18823.1 32895.2 5128.8
0.780 4.725 0.278
3.48 1 2.63
Fresh PIP tin ball
Inactivated one
50
Table 2-4. FT-IR bands assigned to the spectra of precipitates and reference values of Sn(II)-citrate chelate, SnO2・nH2O and Sn(OH)2. FT-IR band position (cm-1) Tselesh (2008) Assignment νOH νHOH
Sn(Ⅱ)-citrate (54.3％ Sn)
SnO2・nH2O
This work Sn(OH)2
3593 3470-3370
3550-3200
3390-3240
Precipitates
Citric acid monohydrate
3501
3500
3398
3417
νC=O
1723
νC-O
1228
νCH(CH2)
2912, 2853
2921, 2851
νasCOO-
1545
1623
1619
νsCOO-
1390
1384.7
1395
δC-OH
1136
1136
1158
δSn-OH terminal
1247
1245
1248
1245
672, 573
658, 533
546
659, 549
δSn-O
51
Table 2-5. The Langmuir-Hinshelwood kinetic parameters. The pH was adjusted to 3 using citric acid or hydrochloric acid. PIP tin ball : 8g, The volume of reaction solution: 2.2 ml With citrate kapp [mM/min] Km [mM-1] Acid yellow 23 Acid red 27 Reactive black 5 Orange II
0.09 0.36 0.04 0.52
0.74 0.26 1.38 0.25
52
With hydrochloric acid kapp [mM/min] Km [mM-1]
Not determined Not determined 0.02 2.12 0.03 4.20
HO +
Na-O3S
HO +
Na-O3S
N=N
SO3-Na+
N=N SO3-Na+
OrangeⅡ
Acid red 27
HO +
Na-O3S
N
N=N
N
HOOC
SO3-Na+
Acid yellow 23
OH NH2 +
-
Na O3SOCH2CH2O2S
N=N
N=N +
Na-O3S
Fig. 2-1.
SO3-Na+
Structure of azo compounds used.
53
SO2CH2CH2OSO3-Na+
Reactive black 5
4mm
Surface SnOX layer Al2O3
Sn
2μm
Fig. 2-2. The structure of PIP tin ball. The SnOx exhibits the surface tin oxide where the stoichiometric ratio ( x ) is determined as 0≦x≦2 by XPS analysis.
54
3
70
2.9
60
2.8
50
2.7
40 2.6
30
pH
Decolorization ratio (％）
80
2.5
20
10
2.4
0
2.3
0
2.5
5 7.5 10 12.5 15 Citric acid concentration (mM)
17.5
Fig. 2-3. Decolorization ratios of Orange II by PIP-tin-ball in 7min under static condition and pHs at various citric acid concentrations. ◆: Decolorization ratio, □: pH, PIP-tin-ball : 8g, Initial concentration of Orange II : 0.3mM, The volume of reaction solution: 2.2ml
55
(I)
(II)
20.6 min
2.1 min
(III)
30.1 min
1A2N
p-ABS
Orange II
Time (min)
Fig. 2-4. HPLC chart of the decolorized products of Orange II in ten minutes reaction. PIP tin ball : 8g, Initial concentration of Orange II : 0.3mM The volume of reaction solution: 2.2ml
56
Rel.Abundance
Citric acid
m/z
Fig. 2-5.
LC-MS-spectrum of the peak at retention time of 2.1min in HPLC analysis.
57
HO HO3S
N=N Orange II
4H+ + 4e-
HO HO3S
NH2
H2N
ABS
Fig. 2-6.
1A2N
Reductive cleavage of Orange II by PIP tin ball
58
0.35
Concentrations (mM)
0.3 0.25 0.2 0.15 0.1 0.05 0 0
20
40
60
80
100
120
Time (min)
Fig. 2-7. Time course of Orange II degradation products by PIP tin ball under static condition. PIP tin ball : 8g; reaction solution volume : 2.2ml; ○:Orange II; ■:ABS; △:1A2N.
59
3 2.5
ln (Co /C)
2 1.5 1 0.5 0 0
10
20
30
40
50
60
Time (min)
Fig. 2-8. Plot of ln(Co/C) versus t during the decolorization of Orange II (Co is the initial concentration and C is the concentration at time„t‟). Air bubbling condition (△), no gas supply condition (◆) and N2 bubbling condition (■).
60
×１０２
250
unused
Intensity (cps)
200
unused inactivated
150
inactivated 100
（b)
（a)
50
504
500
496
492
488
484
480
Binding energy (eV)
Fig. 2-9.
XPS Sn3d spectra of the unused PIP tin ball and the inactivated PIP tin ball.
61
Chapter 3 Comparative study on the decolorization of Orange II by zero-valence tin in citric acid and in hydrochloric acid solutions
3.1. Introduction Zero-valence metals (ZVM) such as iron, zinc and tin are moderately strong reducing agents that are capable of reducing many environmental contaminants [1]. Zero-valence iron has been extensively studied with respect to the decolorization of azo dye [2-6]. However, zero-valence iron has some drawbacks in practical applications. Once Fe0 is in contact with air, its reducing reactivity is reduced [7]. Color removal was decelerated in the alkaline region of pH because the precipitate of ferrous hydroxide on the surface of iron occupied the reactive sites, leading to terminating the reaction [8]. Few research on the decolorization of azo dyes by zero-valence tin has been conducted. In the previous chapter, PIP tin ball performed efficient decolorization of azo dyes in the presence of citrate [9] and it was suggested that the metallic tin component in the structure of PIP tin ball played a major role in the decolorization reaction, and citric acid enhanced the reducing power of zero-valence tin [9]. The object of this chapter was detailed study on zero-valence tin in the decolorization reaction of Orange II both in citric acid and in hydrochloric acid.
3.2. Methods 3.2.1. Materials An azo dye, Orange II
(Tokyo Kasei Kogyo Co., Ltd.) was used without further
purification. Zero-valence tin (trade name:Tin Drops, average density:7.2865 g/cm3, 62
BET surface area:0.0025 m2/g), p-aminobenzene sulfonic acid (ABS) and citric acid were purchased from Kanto Chemical Co., Ltd., and 1-amino-2-naphthol (1A2N) was purchased from Wako Pure Chemical Industries Ltd. These chemical reagents were of guaranteed reagent grade. Tin rod (diameter:6 mm, length:160 mm) was purchased from Alfer Aesar Co., Ltd.
3.2.2. Decolorization procedure by zero-valence tin EVA (ethylene vinyl acetate copolymer) mesh bags containing 7 grams of zero-valence tin were hung in 100 ml beakers. Dye solutions in citric acid or in hydrochloric acid were added in these beakers to start decolorization reaction. All experiments were performed under stirring condition by magnetic stirrer at room temperature in ambient air. In repeated decolorization of Orange II by zero-valence tin, the first reaction was started as described above and then the reaction solution was replaced with a fresh dye solution every fifty minutes to resume the reaction. This procedure was repeated five times.
3.2.3. Analysis Dye concentration of Orange II was determined by measuring maximum absorbance of the dye solution (485 nm) with a spectrophotometer (UV-220 Shimadzu). To identify decolorization products of Orange II by tin, HPLC analysis was performed with Jasco PU-980 delivery system equipped with a UV-VIS detector (875UV Jasco), column oven (TU-100 Jasco) and reverse phase C18 column (AG 120, 0.46×15 cm; Shiseido). Elution condition was as follows; first step; isocratic condition with 0.1 M phosphate buffer (pH 7) for 5 min, second step; 10 min linear gradient of 0.1
63
M phosphate buffer (pH 7) and water-methanol (1:1 v/v) at a flow rate of 1.0 ml/min, final step; isocratic with water-methanol (1:1 v/v) for 10 min. Eluents were detected at 254 nm. To determine open-circuit potentials (OCP) of tin in citric acid solution and in hydrochloric acid solutions, tin rod samples were used as a working electrode and Ag/AgCl saturated electrode was a reference electrode. These electrodes were set in two acid solutions and the potentials were measured using a voltammeter (Model M-1082, Able). The measurement of BET surface area of tin drops was conducted by Shimadzu Co., Ltd.
3.3. Results 3.3.1. Effect of acid concentrations on decolorization of Orange II Zero-valence tin decolorized Orange II in the presence of 5 mM citric acid. HPLC of treated solution of Orange II by tin showed two new peaks attributed to two products. The two products were identified as p-aminobenzene sulfonic acid (ABS) and 1-amino-2-naphthol (1A2N) by comparing the products with authentic samples [9]. These aromatic amines were produced through the reductive cleavage of azo bond of Orange II as shown in Fig. 3-1. Time courses of the concentrations of Orange II and the products of ABS and 1A2N are shown in Fig. 3-2. It was confirmed that decolorization reaction of Orange II by tin produced ABS stoichiometrically. However, the concentration of 1A2N produced was lower than the stoichiometrically calculated value, mainly because of loss of 1A2N by the autooxidation of 1A2N. The autooxidized products of 1A2N were not confirmed because they were not detected by HPLC [3].
64
Figure 3-3 (a) shows the decolorization of 0.1 mM Orange II by tin in 2.5 mM citric acid and in 2.5 mM hydrochloric acid. Tin/citric acid system showed significantly higher decolorization rates of Orange II than tin/hydrochloric acid system. The time courses of open-circuit potentials (OCP) of tin during the experiments in Fig. 3-3 (a) are shown in Fig. 3-3 (b). OCP is generated from the dissolution reaction of tin as follows. Sn → Sn2+＋2e-
(3-1)
The higher negative value of OCP was obtained in citric acid compared with that in hydrochloric acid. At initial concentration of Orange II of 0.1 mM, decolorization reactions were expressed by the pseudo-first-order reaction ( eqn.(3-2) ) －dC/dt=kobsC
(3-2)
,where C is the Orange II concentration (mM), and kobs is the pseudo-first-order reaction rate constant. Integration of equation (3-2) gives ln(Co/Ct)=kobst
(3-3)
,where Co (mM) is the initial concentration of Orange II, and Ct (mM) is the Orange II concentration at a reaction time t (min). A slope of straight line by ln(Co/Ct) against t, gives kobs. Table 3-1 shows the results of kobs calculated from eqn. (3-3) and the open-circuit potential (OCP) after 18 min from the start of experiments in different citric and hydrochloric acid concentrations. When hydrochloric acid was increased from 2.5 mM to 5 mM, the kobs values increased by more than ten-fold and OCP value in 5 mM hydrochloric acid was negatively higher by 40 mV than that in 2.5 mM hydrochloric acid. The kobs values and OCP values were almost the same in different 65
citric acid concentrations. The kobs values in tin/citric acid in 2.5 and 5 mM were larger than those in tin/hydrochloric acid and OCP values in tin/citric acid were negatively higher by more than 100 mV than those in tin/hydrochloric acid.
In 10 mM, the
significantly high value of kobs was obtained in hydrochloric acid, but kobs value of citric acid was similar to those in 2.5 and 5 mM. Repeated decolorization of Orange II was carried out in 2.5 mM citric acid and 10 mM hydrochloric acid. Decolorization ratios and pseudo-first-order reaction rate constants in each of 50-min repeated reaction are shown in Table 3-2. Decolorization ratio was calculated by the following equation. Decolorization ratio (％)=(Co-C50min)/Co×100
(3-4)
,where Co is the initial concentration and C50min is the concentration of Orange II after 50 min at each cycle. Decolorization ratios and pseudo-first-order reaction rate constants in citric acid were not affected by repeated reactions, but the reaction rate constant in the fifth cycle in hydrochloric acid was half of the value in the first decolorization.
3.3.2. Effect of initial dye concentrations on Orange II decolorization Plotting ln(Co /Ct) against t of decolorization reaction at various initial Orange II concentrations (0.1-0.4 mM) in 5 mM citric acid is shown in Fig. 3-4 (a). Table 3-3 shows the first-order reaction rate constants and the initial reaction rate (vo) calculated by the following equation. vo =(-dC/dt)t=0 =kobsCo
(3-5)
The reactions obeyed first-order reaction model in these dye concentrations, and slight decrease in the values of reaction rate constant was found. For the surface processes the
66
reaction rate could be described by the Langmuir-Hinshelwood model (L-H model) which is expressed in the following equation [10]. 1/vo =1/kapp+1/(kappKmCo)
(3-6)
,where kapp and Km are apparent reaction rate constant and adsorption equilibrium constant, respectively. The reciprocal plots of initial reaction rate and initial dye concentration were linear as shown in Fig. 3-4 (b). Heredia et al. reported that the reciprocals of first-order reaction rate constants were proportional to initial concentrations in L-H model [11]. The similar result was also obtained in this study as shown in Fig. 3-5. Figure 3-6 (a) shows the plotting of ln(Co /Ct) against t at various initial Orange II concentrations (0.1-0.4 mM) in 5 mM hydrochloric acid. Although straight line was attained at Co of 0.1 mM, the plots in the range of 0.2-0.4 mM did not give straight lines. Thus, I applied the following relaxation first-order kinetic in which the rate is expressed by first-order reaction against the difference between the concentration at time t, C t and the concentration at equilibrium state, Ce [12].
－dC/dt=kobs(C-Ce)
(3-7)
ln{(Co-Ce)/(Ct-Ce)}=kobst
(3-8)
The model explains that the reaction proceeds toward the equilibrium state in reversible reaction (see Discussion). The equilibrium concentrations Ce were estimated using Guggenheim method as follows. (Ct-Ce) /(Ct+T-Ce) = e kT Ct=Ce(1- e
kT
) + e
(3-9)
kT
Ct+T
(3-10)
,where T is a constant interval time. The value of 10 min was arbitrarily chosen in this study.
By plotting Ct against Ct+T, the straight lines for 0.2-0.4 mM initial 67
concentrations were obtained as shown in Fig. 3-7. The values of Ce estimated from the intercepts and the slopes of the lines of equation (3-10) is shown in Table 3-4 for each Co. Figure 3-6 (b) shows the plotting of ln((Co-Ce)/(Ct-Ce)) versus t. The first-order-reaction rate constants determined from the slopes of the lines and the initial reaction rate estimated by the following equation are shown in Table 3-4. vo= (-dC/dt)t=0 = kobs (Co-Ce)
(3-11)
Initial reaction rates were independent of initial concentrations of azo dye in the range of 0.2 mM-0.4 mM.
3.3.3. Dependence of pH in citrate buffer on the decolorization of Orange II The effect of pH on the decolorization of 0.1 mM Orange II was investigated in 5 mM citrate buffer. Citrate buffers (pH 3-5.5) were prepared by mixing 5 mM citric acid and 5 mM trisodium citrate. The decolorization kinetics at various pHs are shown in Fig. 3-8. The reaction patters were almost the same and the first-order-reaction model was applicable in pH less than 4. However, the decolorization rates significantly decreased at pH 4.5 and almost no reaction occurred at pH 5.
3.4. Discussion Zero-valence tin/citrate system decolorized efficiently Orange II by the reductive cleavage of azo bond.
The BET surface area normalized rate constant for tin/citrate
solution system was calculated as kSA=0.25 lm-2 min-1 (citrate buffer 5 mM, pH3). It was reported that the BET surface area normalized rate constant for decolorization of Orange II by zero-valence iron was kSA=0.18 lm-2 min-1 at pH3 [10]. Although iron has higher standard reducing potential (Fe2++2e- =Fe : -0.44 V) than tin (Sn2++2e- =Sn :
68
-0.138 V), tin/citrate system showed higher reaction rate constant than that of iron in this experiment. This suggests that citrate played an important role in the enhanced decolorization reaction of azo dye. The dissolution of tin is coupled to the reduction of azo dye as an electron acceptor as shown in Fig. 3-9 and Orange II is decolorized by electron generated from dissolution of tin. In hydrochloric acid and citric acid, the reaction rates depended on acid concentrations as shown in Table 3-1 and Fig. 3-8, respectively. It was reported that the dissolution rate of tin was depended on pH [13]. The pHs of citric acid were higher than those of hydrochloric acid under the same acid concentrations, but the pseudo-first-order reaction rate constants in citric acid were larger than those in hydrochloric acid. The open-circuit potentials (OCP) in citric acid were negatively higer by 100 mV than in hydrochloric acid. This suggests that the reducing power of tin in citric acid was stronger than that in hydrochloric acid, leading to the enhancement of electron transfer to Orange II in citric acid. The comparative kinetics described in the section 3.3.2 allows us to propose the different pathways in the reactions in citric and hydrochloric acids. In the reduction of azo dyes, aromatic amines were produced through hydrazo intermediates [3, 7, 14, 15] and thus the reductive cleavage of azo dyes was speculated to proceed through two steps as follows. R1-N=N-R2＋2H++2e-
R1-NH-NH-R2
(3-12)
R1-NH-NH-R2+2H++2e-
R1-NH2+R2-NH2
(3-13)
When the initial concentration of Orange II was more than 0.2 mM in hydrochloric acid, the reaction kinetics were expressed by the equation (3-7) which introduced the
69
equilibrium concentration. Figure 3-10 shows a scheme of Orange II decolorization in hydrochloric acid by introducing the reversible reaction (3-12) where hydrazo intermediate returned to the parent compound. In the higher concentrations of Orange II, the diffusion of the produced hydrazo intermediate into solution would be significant and the hydrazo intermediate is readily reoxidized to azo compound by the oxidants like oxygen [16]. Thus, the further reduction of hydrazo intermediate by tin surface is retarded and pseudo-first-order reaction kinetic was not applicable. Other possibilities for deviation from first-order kinetic are that dissolved Sn2+ reduces the hydrazo intermediate in solution as shown in bracket in Fig. 3-10 or that the diffusion of Sn2+ into solution may be prohibited by adsorbed dye on the tin and thus the reduction of hydrazo intermediate would be reduced in higher dye concentrations. On the other hand, decolorization reaction in citric acid obeyed a first-order kinetic over the examined range of dye concentrations. In addition, Langmuir-Hinshellwood model (L-H model), which explained that the overall rate was controlled by solid surface reaction, was applicable. This indicates the reduction steps of (3-12) and (3-13) were consecutively and rapidly occurred on tin surface. Dissolution reaction of tin in citric acid (H2Cit-) was accompanied by the following reactions [17]. Sn →
Sn2+
+
Sn2+ +H2CitSnCit-+H+
2e-
(3-1)
SnCit-+2H+
(3-14)
SnHCit
SnHCit + 2H2O
(3-15)
SnO2+H2Cit-+2e-+3H+
(3-16)
Dissolved ion of Sn2+ in citric acid formed a complex of Sn(II)-citrate (SnCit-). The protonated complex (SnHCit) adsorbed on the surface of tin was oxidized to SnO2 and four electrons were liberated during the reactions. The generated electrons would be
70
rapidly transferred to azo dye, leading to rapid reduction to aromatic amines. The scheme described above is summarized in Fig. 3-11. In our previous study on the decolorization of azo dyes by tin-immobilised material, the adsorption of a complex Sn(II)-citrate and the formation of SnO2 on tin surface were confirmed [9]. The dissolved Sn2+ is easily oxidized to Sn4+ and Sn4+ is hydrolyzed in acid media to form the highly insoluble Sn(OH)4 and SnO2, which were precipitated on the surface of tin to form oxide film as shown in Fig. 3-9 [18]. The formation of oxide film blocks the active site of tin, leading to reduced reaction. In hydrochloric acid, half-decrease in reaction rate constant was observed in 5-th repeated use and this was due to the oxide film formation to inhibit the decolorization [19]. However, decolorization was successfully repeated in almost the same rate in citric acid as shown in Table 5-2. It is known that the oxide film formed in citrate solution was thin and conductive for electron transfer [20-22]. Therefore, the tin oxide film in citric acid was not so influential to disturb dissolution of tin and decolorization activity of tin. Open circuit potential (OCP) in hydrochloric acid solution was influenced by acid concentrations as shown in Table 3-1. The similar trend was reported on the decolorization of Orange II by zero-valence iron in which the OCP showed a highly negative value at lower pH and the potential was closely related with the efficiency of decolorization of Orange II by iron [10]. When the concentration of hydrogen ion which is primary electron acceptor increased, dissolution of zero-valence metal was thermodynamically feasible [6]. The OCP was highly negative value in citric acid than hydrochloric acid. Gouda et al. [13] reported that OCP of zero-valence tin in the presence of organic acid was related to the stability of chelate between dissolved Sn2＋ and organic acid and then was described by the Nernst equation [13].
71
ESn/Sn2+=EoSn/Sn2++0.0259 log [Sn2+]
(3-17)
, where ESn/Sn2+ is open-circuit potential and EoSn/Sn2+ is standard reduction potential. OCP determined by equation (3-17) is a function of the mole fraction of the free Sn+2 (uncomplexed Sn+2) in solution. In the presence of citrate, the formation of stable complex between tin and citrate considerably reduced the mole fraction of the free Sn+2 in solution. Thus, shifting towards negative values of OCP indicates the increase in the energy of electrons generated from tin dissolution and the stronger reducing power of the electrons. This electrochemical behavior could be related to the enhancement of the reduction of Orange II in citric acid.
72
References
[1]Choi J.-H. and Kim Y.-H., Reduction of 2,4,6-trichlorophenol with zero-valent zinc and catalyzed zinc, J. Hazard. Mater., 166: 984-991 (2009) [2]Cao J., Wei L., Huang Q., Wange L. and Han S., Reduction degradation of azo dye by zero-valent iron in aqueous solution, Chemosphere, 38: 565-571 (1999) [3]Nam S. and Tratnyek P. G., Reduction of azo dyes with zero-valent iron, Water Res., 34: 1837-1845 (2000) [4]Feng W., Nansheng D. and Helin H., Degradation mechanism of azo dye C. I. Reactive red 2 by iron powder reduction and photooxidation in aqueous solutions, Chemosphere, 41: 1233-1238 (2000) [5]Mu Y., Zhang S.-J. and Zhen J.-C., Kinetics of reductive degradation of Orange II in aqueous solution by zero-valent iron, J. Chem. Technol. Biotechnol., 79: 1429-1431 (2004) [6]Zhang H., Duan L., Zhang Y. and Wu F., The use of ultrasound to enhance the decolorization of the C.I. Acid Orange 7 by zero-valent iron, Dyes Pigments, 65: 39-43 (2005) [7]Cheng S.-F. and Wu S.-C., The enhancemet methods for the degradation of TCE by zero-valent metals, Chemosphere, 41: 1263-1270 (2003) [8]Chang M.-C., Shu H.-Y., Yu H.-H. and Sung Y.-Y., Reductive decolorization and total organic carbon reduction of the diazo dye CI Acid black 24 by zero-valent iron powder, J. Chem. Technol. Biotechnol., 81:1259-1266 (2006) [9]Nishide S., Hirai M. and Shoda M., Decolorization of azo dyes by PIP tin balls in citric acid solution, J. Wat. Environ. Technol., 8:85-98 (2010) [10]Mielczarsky J. A., Montes Atenas G. and Mielczarski E., Role of iron surface oxidation layers in decomposition of azo-dye water pollutants in weak acidic solutions, Appl. Catal. B: Environ., 56: 289-303 (2005) [11]Heredia J. B. D., Torregrosa J. Dominguez J. R. and Peres J. A., Oxidation of p-hydoroxybenzoic acid by UV radiation and by TiO2/UV radiation: comparison and modeling of reaction kinetic, J. Hazard. Mater., B83: 255-264 (2001) [12]Hemalatha M. R. K. and Noorbatcha I., An undergraduate physical chemistry experiment on the analysis of first-order kinetic data, J. Chem. Edu., 74: 972-974 (1997) [13]Gouda V. K., Rizkalla E. N., Abd-el-wahab S and Ibrahim E.M., Corrosion behavior in organic acid solutions-I. Tin electrode, Corro. Sci., 21: 1-15 (1981) 73
[14]Zbaida S., Stoddart A. M. and Levine W. G., Studies on the mechanism of reduction of reduction of azo dye carcinogens by rat liver microsomal cytochrome P-450, Chem. Biol. Interactions, 69: 61-71 (1989) [15]Chang J.-H., Chou C., Lin Y.-C., Lin P.-J., and Ho J.-Y., Kinetic characterization of bacterial azo-dye decolorization by Pseudomonas luteola, Wat. Res., 13: 2841-2850 (2001) [16]Zimmermann T., Kulla H. G. and Leisinger T., Properties of purified Orange II azoreductase, the enzyme initiating azo dye degradation by Pseudomonas KF46, Eur. J. Biochem., 129 :197-203 (1982) [17]Almeida C. M. V. B. and Giannetti B. F., Protective film growth on tin in perchlorate and citric acid electrolytes, Mat. Chem. Phys., 69: 261-266 (2001) [18]Hassan H. H. and Fahmy K., Pitting corrosion of tin by acetate anion in acidic media, Int. J. Electrochem. Sci., 3: 29-43 (2008) [19]Stirrup B. N. and Hampson N. A., The active dissolution of tin in acidic chloride electrolyte solutions – a rotating disc study, J. Appl. Electrochem., 6: 353-360 (1976) [20]Giannetti B. F., Sumodjo P. T. A. and Rabockai T., Electrochemical studies with tin electrodes in citric acid solutions, J. Appl. Electrochem., 20: 672-676 (1990) [21] Séruga M. and Metikoš-Huković M., Passivation of tin in citrate buffer solutions, J. Electro. Chem., 334: 223-240 (1992) [22]Abdel Rehim S. S., Sayyah S. M. and El Deeb M. M., Corrosion of tin in citric acid solution and the effect of some inorganic anions, Mat. Chem. Phys., 80: 696-703 (2003)
74
Table 3-1. Pseud-first-order reaction rate constants (kobs) of decolorization of 0.1 mM Orange II and open-circuit potential (OCP) in different acid concentrations.
Acids
Acid conc. [mM]
pH
kobs [min-1]
OCP [mV]
Hydrochloric acid
2.5
2.68
2.1×10-3
-328
5
2.35
2.8×10-2
-370
10
2.11
7.6×10-2
No measurement
2.5
2.95
5.4×10-2
-464
5
2.80
5.5×10-2
-470
10
2.70
5.8×10-2
No measurement
Citric acid
75
Table 3-2. Decolorization ratios and pseudo-first-order reaction rate constants (kobs) at each 50 min cycle in 5 repeated decolorizations in 2.5 mM citric acid and in 10 mM hydrochloric acid (initial concentration of Orange II 0.1mM). Decolorization ratio
kobs
[％]
[min-1]
Hydrochloric acid
Citric acid
Hydrochloric acid
Citric acid
10 mM
2.5 mM
10 mM
2.5 mM
1
98
93
7.6×10-2
5.7×10-2
2
92
92
5.1×10-2
5.0×10-2
3
87
94
4.0×10-2
5.1×10-2
4
86
92
4.0×10-2
5.3×10-2
5
86
92
3.8×10-2
5.3×10-2
Cycle
76
Table 3-3. Pseudo first-order reaction rate constants (kobs) and initial reaction rates (vo) estimated from first-order reaction profile in different initial Orange II concentrations (Co) in 5 mM citric acid. Co
kobs
vo
[mM]
[min-1]
[mM min-1]
0.1
5.4×10-2
5.3×10-3
0.2
4.8×10-2
9.3×10-3
0.3
4.7×10-2
1.4×10-2
0.4
4.3×10-2
1.7×10-2
77
Table 3-4. Decolorization reaction rate constants (kobs), and equilibrium concentration (Ce) from Guggenheim method in different initial Orange II concentrations in 5mM hydrochloric acid.
Co
Ce
kobs
vo
[mM]
[mM]
[min-1]
[ｍM min-1]
0.2
8.1×10-2
6.4×10-2
7.3×10-3
0.3
1.9×10-1
7.2×10-2
7.5×10-3
0.4
3.1×10-1
7.6×10-2
7.5×10-3
78
HO HO3S
N=N Orange II
4H+ + 4e-
HO HO3S
NH2
H2N
ABS
Fig. 3-1.
1A2N
Reductive cleavage of Orange II by tin
79
0.3
Concentration (mM)
0.25 0.2
0.15 0.1
0.05 0
0
20
40
60 80 Time (min)
100
120
140
Fig. 3-2. Time course of degradation of Orange II by tin/5mM citric acid aqueous solution. △;Orange II, ◆;p-aminobenzene sulfonic acid, ■;1-amino-2-naphthol
80
0.1
OrangeⅡconcentration （ｍM）
0.09 0.08 0.07
(a)
0.06 0.05 0.04 0.03 0.02 0.01 0 0
10
20
30
40
50
60
Time (min)
0
Potential (mV vs Ag/AgCl)
-50 -100
(b)
-150 -200 -250 -300 -350 -400 -450 -500 0
10
20
30
Time (min)
Fig.3-3. Decolorization of Orange II in 2.5 mM citric acid (■) and in 2.5 mM hydrochloric acid ( □ ). Time course of the concentration of Orange II (a) and open-circuit potential (b). Initial Orange II: 0.1 mM
81
6
5
ln (Co/C t)
4 3
(a)
2 1 0
0
40
80
120
Time (min) 250
1/vo (min/mM)
200
150
(b)
100
50
0 0
2
4 6 8 1/Co (1/mM)
10
12
Fig. 3-4. The plots of ln(Co/Ct) against time at various initial Orange II concentrations Co(■;0.1, □;0.2 ▲; 0.3, △;0.4mM) (a), and reciplocal plot of initial reaction rate 1/vo against 1/Co (b) in tin/5mM citric acid system.
82
25
1/kobs (min)
20
15
10
5
0 0
0.1
0.2
0.3
0.4
0.5
Co (mM)
Fig. 3-5. The plots of 1/kobs against Co in tin/ 5 mM citric acid system.
83
1.6 1.4
ln(Co/Ct)
1.2
(a)
1
0.8 0.6
0.4 0.2
0 0
20
40
60
40
60
Time (min)
3.5
ln ((Co-Ce)/(C t-Ce))
3
(b)
2.5
2 1.5
1 0.5
0 0
20
Time (min)
Fig. 3-6. The plots of ln(Co/Ct) against time at various initial Orange II concentrations Co(■ ;0.1, □;0.2, ▲ ; 0.3, △;0.4 mM) (a), and the plots of ln((Co-Ce)/(Ct-Ce)) against time (b) in tin/5 mM hydrochloric acid system.
84
0.25
C t (mM)
0.2 0.15 0.1
(a)
0.05 0 0.07
0.09
0.11
0.13
0.15
Ct+T (mM)
0.35
0.45
0.3
0.4 0.35 0.3 C t (mM)
C t (mM)
0.25 0.2 0.15 0.1
0.25 0.2 0.15
(b)
(c)
0.1
0.05
0.05
0
0 0.19
0.21
0.23
0.25
0.27
0.3
Ct+T (mM)
0.32
0.34 Ct+T (mM)
0.36
0.38
Fig. 3-7. The plots of Ct against Ct+T at various initial OrangeII concentrations (a); 0.2, (b); 0.3, (c); 0.4 mM during decolorization reaction with tin/5mM hydrochloric acid system.
85
OrangeⅡ concentration (mM)
0.12 0.1 0.08 0.06
0.04 0.02 0
0
40
80
120 160 200 Time (min)
240
280
Fig. 3-8. The time course of the decolorization of Orange II at various pH values in tin/citrate buffer system. ○;pH3, ◆;pH3.5, ◇;pH4, □;pH4.5, ▲;pH5, △;pH5.5
86
Sn
Dissoluton
AZO
2e-
Sn+2 Oxidation
Sn+4 Sn(OH)4
SnO2
Hydrolysis
Precipitation Dehydration
Fig. 3-9. The scheme of the dissolution of tin in acid media.
87
Oxide film formation (passivation)
Surface
Solution Diffusion
Orange II
Sn Dissolution
Orange II
2e-
Sn+2 Diffusion hydrazo
hydrazo
Sn+2
2e2eDissolution
Sn+2 ABS
ABS
1A2N
Sn+4 1A2N
Diffusion
Fig. 3-10. The scheme of the decolorization of Orange II in hydrochloric acid. The reduction by stannous ion in solution is shown in brackets.
88
Surface
Sn
Solution Diffusion Orange II
Orange II
4e-
Diffusion ABS
1A2N
ABS
Sn(II)-Cit Dissolution (Sn(II)-HCit)ad Adsorption Oxidation SnO2
Fig. 3-11. The scheme of the decolorization of Orange II in citric acid.
89
1A2N
Chapter 4 Biodegradation of two aromatic amines produced from the decolorization of Orange II by zero-valence tin
4.1. Introduction Azo dyes can be decolorized by bacteria under anaerobic conditions, with the production of aromatic amines. Only a few simple aromatic amines are degraded under anaerobic conditions [1, 2]. As aromatic amines are more easily degraded under aerobic conditions than anaerobic conditions, mineralization of an azo dye using an alternating anaerobic-aerobic process [3], has attracted attention as a practical method [4-6]. However, anaerobic process often requires long reaction time and consequent removal of aromatic amines under aerobic conditions is affected by the previous process [6]. On the other hand, decolorization of azo dye by chemical reduction is rapid, and this reduction can be a possible pretreatment for subsequent aerobic biodegradation. Iron-treated dye solutions exert a higher biochemical oxygen demand (BOD) than the solutions containing the untreated dye solution mainly because azo dyes are more recalcitrant than the iron-treated dye solutions [7]. Thus, zero-valence metal treatment of azo dyes will be effective to increase the efficiency in subsequent aerobic biological treatments. In the previous chapter, it was found that p-aminobenzene sulfonic acid (ABS) and 1-amino-2-naphthol (1A2N) were generated after tin-treatment of Orange II. In this chapter, the purpose was the biological treatment of the mixture of ABS and 1A2N produced after tin-treatment of Orange II. ABS is discarded to the rivers and surface waters due to its high water solubility [8] and 1A2N is unstable and highly toxic [9, 10]. Since in the chemical decolorization of the azo dye, Orange II using zero-valence tin
90
does not reduce the carbon numbers of the azo dye, I expected that subsequent biological treatment of aromatic amines will be effective to reduce the carbon number of these amines. For the establishment of effective biological treatment, I conducted enrichment culture to find effective microorganisms to degrade two aromatic amines. In addition, co-culture of the isolated strains to degrade the mixture of two aromatic amines produced from tin-treatment of Orange II was conducted in the absence and in the presence of tin.
4.2. Methods 4.2.1. Materials Orange II was purchased from Tokyo Kasei Kogyo Co. Ltd. (Tokyo); p-aminobenzene sulfonic acid (ABS) and tin (trade name: Tin drops) were purchased from Kanto Chemical Co., Ltd. (Tokyo) and 1-amino-2-naphthol (1A2N) was purchased from Wako Pure Chemical Industries Ltd.(Tokyo). These chemical reagents were of guaranteed reagent grade.
4.2.2. Enrichment culture Activated sludge from the local sewage treatment plant in Machida (Tokyo) was used as a seed for the three enrichment cultures. In the cultures, runs 1 and 2, 200 mg/l ABS and 200 mg/l 1A2N were used, respectively as the sole carbon source. The culture, run 3 included both 100 mg/l ABS and 100 mg/l 1A2N. Activated sludge of 10 ml was inoculated into three 500 ml shaken flasks containing ABS and/or 1A2N in 100 ml mineral medium. The mineral medium is composed of 1.55 g/l K2HPO4, 0.85 g/l KH2PO4, 2.0 g/l (NH4)2SO4, 0.1 g/l MgCl2・6H2O, and 0.13 vol % trace mineral solution.
91
Trace mineral solution is composed of 10 g/l EDTA, 2 g/l ZnSO4, 1 g/l CaCl2･2H2O, 5 g/l FeSO4･7H2O, 0.2 g/l Na2MoO4.7H2O, 0.2 g/l CuSO4･5H2O, 0.4 g/l CoCl2･6H2O, and 1 g/l MnCl2･4H2O. These flasks were shaken at 30 oC at 120 strokes per min (spm). After 14 days, each culture fluid of 10 ml was inoculated into the each fresh medium and this procedure was repeated. In run 1, the culture was repeated at the time when ABS was completely eliminated. In run 2, the culture was repeated at intervals of 7-10 days. In run 3, 11-th repeated culture was used for further three enrichment cultures, containing 200 mg/l ABS or 200 mg/l 1A2N and a mixture of 100 mg/l ABS and 100 mg/l 1A2N, respectively as shown in Table 4-1.
4.2.3. Isolation of amine-degrading strains The enrichment culture of 0.1 ml was inoculated onto nutrient agar plates and then incubated at 30 oC. The nutrient agar plate contained 18 g/l nutrient broth (Eiken Co., Tokyo) and 15 g/l agar (Taiyo-agar, Shimizushokuhin Co., Shizuoka). After two days, several colonies which showed different morphologies were scraped by platinum wire loop to inoculate them in test tubes containing 10 ml nutrient broth (18 g/l). The test tubes were shaken at 30 oC at 120 spm for 24 hr. The grown culture fluid was diluted with sterilized water and spread onto nutrient agar plate. Single colonies appeared were used to confirm degradation of each amine.
4.2.4. Preservation of isolated strains The cultures of ABS- or 1A2N-degrading strains grown in nutrient broths (18 g/l) were preserved -80 oC in 50 % glycerol at 1:1 (v/v) ratio in microtubes. When each strain in ABS enriched culture which showed ABS degradation was separated and grown in 92
ABS-containing mineral medium, no isolate showed degradation of ABS. Therefore, we preserved the enriched mixed culture as ABS-degrading consortium.
4.2.5. Degradation of aromatic amines by isolated strains A 0.1-ml portion of stock solution of isolated 1A2N-degrading strains or ABS-degrading consortium was inoculated into test tubes containing 10 ml nutrient broth (18 g/l), and then incubated at 30 oC at 120 spm for 24 hr. The grown microorganisms were harvested by centrifugation at 10,000 rpm for 10 min. The supernatant was excluded, and then a concentrated cell pellet was washed with sterilized water. The cell suspension was centrifuged at 10,000 rpm for 10 min and the supernatant was excluded. Then, the cell suspension in 1 ml sterilized water was inoculated into each 500 ml flask containing 200 mg/l ABS or 200 mg/l 1A2N in 100 ml mineral medium.
4.2.6. Preparation of tin-treated solution of Orange II Seven grams of tin were put into an EVA (ethylene vinyl acetate copolymer) meshed bag and the bag was placed in a 500 ml round-bottom flask. The 100 ml dye solution of 0.3 mM Orange II in 20 mM hydrochloric acid (pH=1.7) was added in the flask and stirred for 2 hr by magnetic stirrer for decolorization of Orange II. After reaction, the orange color of the solution was completely disappeared. The decolorized solution was neutralized by 0.2 N sodium hydroxide solution, and was supplemented with the ingredients of mineral medium described in section of 4.2.2.
93
4.2.7. Co-culture treatments of a mixture of authentic samples of ABS and 1A2N or tin-decolorized solution of Orange II P. aeruginosa m2 and ABS-degrading consortium were grown in nutrient broth as described above, and both cell pellets which were obtained by centrifugation and twice washing with sterilized water were inoculated into the mixture of authentic samples of ABS and 1A2N or the 100 ml tin-treated solution of Orange II described above and co-cultures were conducted. In these experiments, the aromatic amines were the carbon sources for the co-cultures.
4.2.8. Analysis The optical density (OD) of culture broth was measured at 660nm using spectrophotometer (UV-1200, Shimazu, Kyoto). To determine the concentration of ABS, HPLC analysis was performed with Jasco PU-980 delivery system equipped with a UV-VIS detector (875UV Jasco, Tokyo), a column oven (TU-100 Jasco, Tokyo) and a reverse phase C18 column (AG 120, 0.46 cm×15 cm; Shiseido, Tokyo). Elution was conducted under isocratic conditions using mixture of 40 % of 0.05 M phosphate buffer (pH 6.5) and 60 % methanol at a flow rate of 1.0 ml/min. Detection of eluted materials was carried out by monitoring absorption at 254 nm. Chemical Oxygen Demand (COD) with chromate was determined by a closed reflux titrimetric method to titrate by standard ferrous ammonium sulfate after decomposition of samples by heating at 150 oC for 2 hr [11].
94
4.3. Results 4.3.1. Enrichment culture In the previous chapter, tin decolorized Orange II by the following reductive cleavage of azo group to produce ABS and 1A2N. HO
HO H3OS
Orange II
N=N
+4H++4e
H3OS
-
NH 2
ABS
+
H2N
1A2N
For subsequent biological degradation of ABS and 1A2N to reduce carbon numbers, enrichment cultures were conducted. In the enrichment culture of run 1 in the mineral medium containing ABS as a carbon source, the degradation of ABS was observed after 8 days and ABS was completely eliminated after 14 days. When this culture broth was inoculated into a fresh medium including 200 mg/l ABS, the complete elimination of ABS was confirmed after 3 days and the increase in OD at the depletion of the peak of ABS in HPLC analysis was observed. Several microorganisms were isolated on the nutrient agar plates. However, no purified strains showed ABS-degrading ability. Therefore, I used the mixed culture as ABS-degrading consortium for the following experiments. As 1A2N was unstable and was converted to a compound which could not be detected by HPLC analysis, the degradation of 1A2N was monitored by the COD removal. In run 2, the COD removal ratio was 40% in 3 days in the second enrichment culture after 14-day first enrichment culture, indicating that 1A2N-degrading microorganisms as the sole carbon source were enriched. In further experiments, the COD removal ratio was almost constant at 40% and one morphological type of round and ocher-colored colony became dominant in the plate assay, which was named as strain AN1. The color of
95
culture medium containing 1A2N as the sole carbon source was brown, presumably because 1A2N was auto-oxidized to naphthoquinone, and then a dimer was formed between naphtoquinone and 1A2N [12]. As the brown color of the culture remained after each repeated culture, this dimerization may have made the further biodegradation of 1A2N difficult. Run 2 was finished at the 8th repeated culture, when COD removal ratio was maintained at 40%. When the removal ratio of COD by the strain AN1 was compared with that by another 1A2N-degrading microorganism, m2, which was isolated in run 3, m2 showed a significantly higher removal ratio and thus, m2 was used for further experiments. In the first enrichment culture of run 3, which contained both ABS and 1A2N, the degradation ratio of ABS measured by HPLC was 17% in 14 days, which was lower than that in run 1. As significant decrease in COD was not observed in the first culture, this suggests no degradation of 1A2N. In the second culture, the complete depletion of ABS was observed at 20 days, and COD removal ratio reached 38%. In the third culture, ABS was completely degraded in 8 days, and COD removal ratio increased to 50%. In the 5th culture, two different colonies (m1, m2) became dominant. The purified two strains showed degradation activity only against 1A2N, and m2 strain showed higher activity. As two strains did not show the degradation activity against ABS, ABS-degrading microorganisms may exist in miner members in this culture. Figure 4-1(a) shows change in ABS concentration and COD in the 10th repeated culture. Figure 4-1(b) shows the cell numbers of m1 and m2 strains. The cell numbers were counted by plating diluted culture broth on the nutrient agar. During 2 days of cultivation, almost no ABS was degraded and the significant growth of m1 and m2 suggested the consumption of 1A2N as growth substrate by m1 and m2, leading to the
96
slight decrease in COD. The cell numbers of m1 and m2 leveled off after 2 days and the decrease in COD after 2 days was reflected by the consumption of ABS. The compound was completely degraded and COD removal was 80% in 6 days. A characteristic peak at retention time of 3.9 min in HPLC chart was constantly found throughout the degradation test, which is assumed to correspond to a complex of ABS and 1A2N through auto-oxidation. The complex was considered to be different from Orange II because Orange II was not detectable under this HPLC analysis condition. The complex showed reddish brown color and remained even after ABS was completely degraded. In the 11th repeated culture, the culture fluid was transferred into three kinds of media, the medium of the sole addition of ABS or 1A2N or the medium containing a mixture of ABS and 1A2N. Table 4-1 shows the ABS concentrations and COD removal ratios at different times in these cultures. In the culture containing mixture of ABS and 1A2N, ABS was completely degraded after 6 days and COD removal was 83%. The reddish brown complex remained even after 6 days. In the sole addition of ABS or 1A2N, both amines were completely degraded after 6 days. After 1A2N was completely degraded, the culture was spread on nutrient agar plate, and m2 strain became dominant. According to 16S rRNA analysis, m2 strain showed 99% homology to P. aeruginosa. P. aeruginosa was reported to show degradation activity against naphthol structure [13]. Since P. aeruginosa m2 as 1A2N-degrader showed higher COD removal than AN1 found in run 2, m2 was used in the succeeding experiments. From the enrichment culture with the sole addition of ABS, several microorganisms were also isolated on the nutrient agar plates. However, similar to run 1, no purified strains showed ABS-degrading ability. As the ABS-degrading consortium found in run 1 showed faster removal rate than that in run 3 (data not shown), I used the mixed culture obtained in
97
run 1 as ABS-degrading consortium for further experiments.
4.3.2. Degradation kinetics of ABS by ABS-degrading consortium Figure 4-2 shows the ABS degradation kinetics by ABS-degrading consortium. To investigate the effect of nitrogen source, different concentrations of (NH4)2SO2 in the range of 0～2.0 g/l were added to the culture. Although the 4-day lag phases which would be associated with the induction time for the degradation of ABS were observed, ABS was depleted within 10 days and ABS degradation kinetics were almost the same at varied concentrations of (NH4)2SO2. The repeated cultures at the concentrations of 0.5 g/l and 0 g/l of (NH4)2SO2 were conducted as shown in Fig. 4-3. Each culture fluid of 10 ml was inoculated into 100 ml of fresh mineral medium containing 200 mg/l ABS. The days to show complete degradation in the second culture at 0 g/l and 0.5 g/l of (NH4)2SO2 were 5 and 3 days, respectively. In the third culture, the complete degradation time was shortened to 2 days at 0.5 g/l (NH4)2SO2 and 3 days at 0 g/l. As the complete degradation was observed even in the absence of (NH4)2SO2, it is assumed that the consortium can use ABS not only as a carbon source but also as a nitrogen source. This indicates a possibility of mineralization of ABS by the ABS-degrading consortium.
4.3.3. Degradation of 1A2N The degradation of 200 mg/l 1A2N by m2 strain is shown in Fig. 4-4. The cell growth was observed along with COD removal and 80% COD removal was attained in one day. Under the same experimental condition, AN1 strain isolated from the enrichment culture run 2 showed at most 40% COD removal as shown in Table 4-2. I compared
98
degradation of 1A2N by several strains isolated in my laboratory which showed degradation activities against aromatic compounds [14, 15]. The COD removal ratios after two days and the initial cell numbers are shown in Table 4-2. No degradation activity against 1A2N was found in strains No. 4 and BTO 61. AN1 strain isolated from the enrichment culture run 2 and NBM 21 strain attained around 40% COD removal ratio. The m2 strain achieved the highest COD removal ratio among these strains. Therefore, the strain m2 was used for further experiments as described below.
4.3.4 Co-culture of ABS-degrading consortium and m2 in the tin-treated solution of Orange II After the complete decolorization of Orange II (0.3 mM) by 7g tin entrapped in a meshed bag for 2 hr, ABS and 1A2N were produced. Tin was then removed, the tin-free solution was neutralized and the mineral medium, ABS-degrading consortium and m2 were added. The time courses of the concentrations of ABS, COD, the cell number of m2 strain and the peak area at retention time, Rt=3.9 min, in HPLC are shown in Fig. 4-5. The degradation of ABS was slight for 5 days. The cell number of m2 increased for 2 days (Fig. 4-5 (c)), indicating that m2 strain grew by utilizing 1A2N for 2 days and then decreased afterward. Thus, 15% of COD removal observed in 2 days was considered to correspond to microbial degradation of 1A2N. The COD removal which included degradation of ABS and 1A2N was 50% in 7 days. When the tin-decolorized solution was neutralized, the color of the solution became reddish brown mainly due to the formation of the complex by auto-oxidation of ABS and 1A2N. In addition, the peak at retention time, Rt=3.9 min in HPLC analysis attributed to the complex was observed and this peak area was maintained during the
99
experimental period (Fig. 4-5(d)), indicating that the initial abiotic formation of the complex was dominant. As the structure of the complex was unknown, I could not determine the concentration of the complex. However, the peak area reflected the existence of the complex. I also conducted the co-culture of ABS-degrading consortium and m2 strain using the mixture of authentic samples of ABS (50 mg/l initial concentration) and 1A2N (50 mg/l initial concentration). Figure 4-6 shows the change in ABS, COD, the number of m2 and the peak area at retention time, Rt=3.9 min. The patterns of change in ABS and COD were similar to those in Fig. 4-5. The change in the cell number of m2 was much clearer than that in Fig. 4-5 mainly because the initial concentration of m2 was lower in Fig. 4-6. The co-culture of ABS-degrading consortium and m2 or sole inoculation of m2 strain was carried out in 50 mg/l of ABS and 100 mg/l of 1A2N. The two values of COD removal ratios at 2 days when ABS degradation did not occur and when complete ABS depletion was observed are shown in Table 4-3. The COD removal at 2 days was mainly attributed to the degradation of 1A2N, because ABS was not significantly degraded at this time. The significant increase in the COD removal at 2 days at 100 mg/l of 1A2N was observed. When only m2 strain was inoculated in the mixture of ABS and 1A2N, no degradation of ABS was observed and the COD removal ratio at 2 days was 44%. This value was similar to the result of co-culture in 50 mg/l of ABS and 100mg/l of 1A2N. This result indicated that m2 strain exerted the degradation activity against 1A2N and the activity in the co-culture was not different from that with single inoculation. The estimation of the activity of ABS-degrading consortium in co-culture was carried
100
out. In Fig. 4-2, the concentration of ABS (220mg/l) was equivalent to COD (330mg/l), and thus the ratio of COD to ABS was 1.5. In Fig. 4-6, after 3 days, the amount of ABS degraded by ABS-degrading consortium was 30mg/l. During this period, the decrease in COD was 45mg/l, leading to a 1.5 ratio of COD to ABS. This also indicates that ABS-degrading consortium exerted its activity in the co-culture which is similar to that in the single inoculation. The COD removal ratios at the complete depletion of ABS were in the range of 50-67% (Table 4-3). The main reason for the incomplete COD removal even after the complete depletion of ABS may be the complex formation between ABS and 1A2N by auto-oxidation which was not degraded by isolated strains (see Fig. 4-5 d and Fig. 4-6 d). When the co-culture of ABS-degrading consortium and m2 was mixed with tin-treated solution of Orange II in the presence of 7 g tin, the degradation of ABS occurred from the start of co-culture as shown in Fig. 4-7 (a) and the peak of HPLC which was attributed to the complex vanished at 2 days as shown in Fig. 4-7 (b). This indicated that the complex was degraded by tin. Along with the degradation of the complex, the reddish brown color of the complex disappeared (data not shown). When the tin-treated solution of Orange II was shaken in the presence of 7 g tin without inoculation of ABS-degrading consortium and m2 at neutral pH, no elimination of ABS and COD was observed (date not shown).
4.4. Discussion On the purpose of the biological degradation of two aromatic amines, ABS and 1A2N which were produced after decoloization of Orange II by tin, the enrichment cultures by
101
activated sludge were performed. I obtained ABS-degarding consortium by the enrichment culture using ABS as a sole carbon source. Aromatic compounds with sulfonic acid group have difficulty in transporting across cell membrane, and thus uptake of these compounds needs induction for the specific transport system [16]. Because ABS has high polar structure with sulfonic acid group, this may be the reason why ABS is recalcitrant for biodegradation. Some single strains were reported to degrade ABS [8, 17-19]. Degradation of ABS by a mixed culture indicated that mutual interactions induced the degradation of ABS [20]. In this study, ABS degradation was observed only by the isolated consortium. A few strains would be related with the degradation of ABS in this study. When ABS degradation kinetics by ABS-degrading consortium were examined, the lag phase was observed in the degradation of ABS. It was suggested that the lag phase would be responsible for the induction time for degradation of ABS. The lag phase was considerably short in the repeated culture because consortium had already induced. Singh et al. reported that 800 mg/l of ABS was completely degraded in 14 hr by the induced culture of Agrobacterium sp [19]. This fast degradation was mainly due to the rich complex medium which included yeast extract as a growth factor. In this study, only synthetic medium was used and thus, the degradation by this consortium will be enhanced in complex medium. Biodegradation of 1A2N was less intensively studied compared with that of ABS. A few reports speculated the degradation of 1A2N by dye-degrading strains after degradation of dyes [21, 22], but no data on the degradation of 1A2N have been shown after single supply of 1A2N. These reports did not clearly show the extent of degradation of 1A2N. As this compound undergoes rapid auto-oxidation, it is difficult to
102
determine quantitative degradation of 1A2N. Therefore, to estimate the extent of biological degradation of 1A2N, COD analysis was used in this study. In the enrichment culture using 1A2N as a carbon source, the COD removal was around 40 %. When AN1 strain which was obtained by this enrichment culture was used as degradation test of 1A2N, the COD removal was 43 %. In the enrichment culture including ABS and 1A2N simultaneously, I could isolate m2 strain (P. aeruginosa) which has high degradation ability against 1A2N. Under the condition including ABS and 1A2N simultaneously, the medium was colored, which may be responsible for the formation of complex of these two amines. It was considered that P. aeruginosa m2 strain which has resistance against the complex could grow predominantly in the enrichment culture including ABS and 1A2N. By co-culture of m2 strain and ABS-degrading consortium with a mixture of ABS and 1A2N, I clarified that ABS-degrading consortium and m2 strain exerted each degradation activity against the mixture of aromatic amines. The similar degradation characteristics by the ABS-degrading consortium and m2 strain were obtained in the tin-treated Orange II solution, although the colored recalcitrant complex, which was formed by auto-oxidation of ABS and 1A2N, remained. In the presence of tin in the co-culture, the complex was eliminated and the degradation of ABS was enhanced. The details of this mechanism are not clear at this moment. The biodegradation of ABS takes part in enzyme system such as deoxygenase [17]. It was reported that the desulfonation in the biodegradation pathway of ABS was occurred by the dioxygenase, and thus the activation of the enzyme needs the coenzyme such as NADH [23].
In the presence of
tin which shows powerful reducing activity, the regeneration of reduced form of NADH would be enhanced, resulting in rapid uptake of ABS.
103
An azo dye was degraded in a relatively short time of 2 hr by tin treatment and aromatic amines that are more amenable to aerobic biodegradation were produced. This result suggests that sequential tin treatment and aerobic biodegradation may be a feasible method to treat wastewaters containing azo dyes. In many effluents in practical textile industries, organic load includes starches and their derivatives and these organic compounds are readily degraded under conventional aerobic conditions [24]. However, persistent dyes would be remaining after the treatment. Under this situation, the two-stage-treatment of chemical reduction by using zero-valence tin and biological oxidation by the co-culture of amine-degrading microorganisms will be applicable without extra carbon sources.
104
References [1]Tan N. C. G., Leeuwen A., Voorthuizen E. M., Slenders P., Prenafeta-Boldu F. X., Temmink H., Letting G. and Field J. A., Fate and biodegradation of sulfonated aromatic amines, Biodegradation, 16: 527-537 (2005) [2]Pandey A., Singh P. and Iyengar L., Bacterial decolorization and degradation of azo dyes, Int. Biodet. Biodeg., 59: 73-84 (2007) [3]Haug W., Shmidt A., Nörtemann B., Hempel D. C., Stoltz A. and Knackmuss H.-J., Mineralization of the sulfonated azo dye mordant yellow 3 by a 6-aminonaphthol -2-sulfonate –degrading bacterial consortium, Appl. Environ. Microbiol., 57: 3144-3149 (1991) [4]Stolz A., Basic and applied aspects in the microbial degradation of azo dyes, Appl. Microbiol. Biotechnol., 56: 69-80 (2001) [5]Pearce C. I., Lloyd J. R. and Guthrie J. T., The removal of color from textile wastewater using whole bacterial cells: a review, Dyes Pigments, 58: 179-196 (2003) [6]Van der Zee F. P. and Villaverde S., Combined anaerobic-aerobic treatment of azo dyes-A short review of bioreactor studies, Water Res., 39: 1425-1440 (2005) [7]Perey J. R., Chiu P. C., Huang C.-P. and Cha D. K., Zero-valent iron pretreatment for enhancing the biodegradability of azo dyes, Water Environ. Res., 74: 221-225 (2002) [8]Wang Y.-Q., Zhang J.-S., Zhou J.-T. and Zhang Z.-P., Biodegradation of 4-aminobenzenesulfonate by a novel Pannonibacter sp. W1 isolated from activated sludge, J. Hazard. Mater., 169: 1163-1167 (2009) [9]Gottlieb A., Shaw C., Smith A., Wheatley A., and Forsythe S., The toxicity of textile reactive azo dyes after hydrolysis and decolorization, J. Biotechnol., 101: 49-56 (2003) [10]Carvalho C., Fernandes A., Lopes A., Pinheiro H. and Gonçalves I., Electrochmical degradation applied to the metabolites of Acid Orange 7 anaerobic biotreatment, Chemosphere, 67: 1316-1324 (2007) [11] APHA., Standard Methods for the Examination of Water and Wastewater. 17 th edn, American Public Health Association/American Water Works Association/Water Pollution Control Federation, Washington DC, USA. (1989) [12]Pearce C. I., Christie R., Boothman C., Constein H. V. Guthrie J. T. and Lloyd J. R., Reactive Azo dye reduction by Shewanella strain J18 143, Biotechnol. Bioeng., 95: 692-703 (2006) [13]Nachiyar C. V. and Rajakumar G. S., Biodegradation of 8-anilino-1-naphthalenesulfonic acid by Pseudomonas aeruginosa, J. Ind. Microbiol. 105
Biotechnol., 33: 845-849 (2006) [14]Joo H.-S., Removal characteristics of high-strength ammonium by Alcaligenes faecalis with heterotrophic nitrification and aerobic denitrification. PhD thesis, Department of Environmental Chemistry and Engineering, Tokyo Institute of Technology, (2005) [15]Jeong E., Charastaristic of xylene degradation by newly isolated bacteria in biofilter. PhD thesis, Department of Environmental Chemistry and Engineering, Tokyo Institute of Technology, (2007) [16]Locher H. H., Poolman B., Cook A. M. and Konings W. N., Uptake of 4-toluene sulphonate by Comamonas testosterone T-2, J. Bacteriol., 175: 1075-1080 (1993) [17]Junker F., Leisinger T. and Cook A. M., 3-sulphocatechol 2,3-dioxygenase and other dioxygenase in the degradative pathways of 2-aminobenzenesulphonic, benezenesulphonic and 4-toluenesulphonic acids in Alcaligenes sp. strain O-1, Microbiology, 140: 1713-1722 (1994) [18]Perei K., Rákhely G., Kiss, I., Polyák B. and Kovács, K. L., Biodegradation of sulfanilic acid by Pseudomonas paucimobilis, Appl. Microbiol. Biotechnol., 55: 101-107 (2001) [19]Singh P., Mishra L. C., Pandey A. and Iyengar L., Degradation of 4-aminobenzenesulfonate by alginate encapsulated cells of Agrobacterium sp. PNS-1, Biores. Technol., 97: 1655-1659 (2006) [20]Fiegel B. J. and Knackmuss H.-J., Syntrophic interactions during degradation of 4-aminobenzenesulfonic acid by a two species bacterial culture, Arch. Microbiol., 159: 124-130 (1993) [21]Coughlin M. F., Kinkle B. K. and Bishop P. L., Degradation of acid orange 7 in an aerobic biofilm, Chemosphere, 46: 11-9 (2002) [22]Xu M., Guo J. and Sun G., Biodegradation of textile azo dye by Shewanella decolorationis S12 under microaerophilic conditions, Appl. Microbiol. Biotechnol., 76: 719-726 (2007) [23]Thurhheneer T., Köheler T., Cook A. M. and Leising T., Orthanilic acid and analogous as carbon sources for bacteria: Growth physiology and enzyme desulphonation, J. Microbiol., 132:1215-1220 (1986) [24] Delée W., O‟Neill C., Hawkes, F. R. and Pinheiro H. M., Anaerobic treatment of textile effluents: a Review, J. Chem. Technol. Biotechnol., 73: 323-335 (1998)
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Table 4-1.
Degradation of ABS and/or 1A2N by 11-th enrichment culture in run 3
Carbon sources
Time (days)
ABS (mg/l)
COD removal ratio (%)
0 3 6
194 176 0
Nt＊ Nt＊ Nt＊
0 3 6
0 0 0
0 100 100
0 3 6
85 56 0
0 50 83
ABS 200 mg/l
1A2N 200 mg/l
ABS 100 mg/l and 1A2N 100 mg/l
＊
Not tested
107
Table 4-2.
Comparison of 1A2N degradation activity COD removal ratio (％) (2day culture time)
Initial cell numbers (cfu/ml)
References
Strains AN1
40
1×106
This study
m2 (Pseudomonas aeruginosa)
80
1.2×108
Pseudomonas sp. NBM21
43
6×107
0
4×109
Joo (2005)
0
8×106
Jeong (2007)
Alcaligenes faecalis No. 4 Rhodococcus sp. BTO61
108
Jeong (2007)
Table 4-3. Degradation of authentic ABS and 1A2N by single culture of m2 or co-culture of ABS-degrading consortium and m2.
Carbon sources (mg/l) ABS 1A2N
50 50 50
100 100 50
COD removal ratio (%) After 2 days After complete ABS depletion (Time)
Inoculation
Consortium* Consortium*
*ABS-degrading consortium
109
m2 m2 m2
44 44 20
― 67 (7 days) 50 (6 days)
80
200 180 160
(a)
140 120
40
100 80
COD (mg/l)
ABS (mg/l)
60
60
20
40 20
0
0 0
1
2
3
4
5
6
7
6
7
Time (days)
6.5
Cell number, log (cfu/ml)
6 5.5 5 4.5
(b)
4 3.5 3 0
1
2
3
4
5
Time (days)
Fig.4-1. Change in ABS concentration (□) and COD (■) during the 10th repeated enrichment culture in run 3 which contained ABS and 1A2N (a) and the time course of cell number of m1(●) and m2 (▲) strains which predominantly grew in the enrichment culture (b).
110
250
ABS (mg/l)
200
150
100
(a)
50
0 0
2
4
6
8
10
12
8
10
12
Time (days) 400 350
COD (mg/l)
300 250 200 150 100
(b)
50 0 0
2
4
6 Time (days)
Fig. 4-2. Changes in the concentration of ABS (a) and COD (b) during the degradation of ABS by ABS-degrading consortium at 0 g/l (◆), 0.5 g/l (◇), 1 g/l (△), and 2 g/l (▲) of (NH4)2SO4.
111
ABS, COD (mg/l)
500 450 400 350 300 250 200 150 100 50 0
0.18 0.16 0.14 0.12 0.1 0.08 0.06 0.04 0.02 0
(a)
0
1
2
3
4
5
6
OD (660nm)
3-rd
2-nd
7
350
0.16
300
0.14
250
0.12
200
0.1
(b)
0.08
150
0.06
100
0.04
50
0.02
0
OD (660nm)
ABS, COD (mg/l)
Time (days)
0 0
1
2
3
4 5 Time (days)
6
7
8
9
Fig. 4-3. Repeat cultures in Fig.2 at 0.5 g/l (NH4)2SO4 (a) and without (NH4)2SO4 (b). The concentrations of ABS (□), COD (■), and OD (○).
112
8.4
160
8.35
COD (mg/l)
140 8.3
120 100
8.25
80
8.2
60
8.15
40
Cell number, log (cfu/ml)
180
8.1
20 0
8.05 0
1
2
3
Time (days)
Fig. 4-4. by m2.
Changes in COD (●), and cell number (○) during the degradation of 1A2N
113
40
160
35
140 120
25
COD (mg/l)
ABS (mg/l)
30
20 15
(a)
10
100
80 60
5
20
0 0
2
4
6
8
0
10
0
Time (days)
9.3
80000
9.2
70000
9.1
60000
9 8.9 8.8 8.7
(c)
8.6
2
4
6
8
10
8
10
Time (days)
Peak area
Cell number, log (cfu/ml)
(b)
40
50000 40000 30000
(d)
20000 10000
8.5
0
8.4 0
2
4
6
8
0
10
Time (days)
2
4
6
Time (days)
Fig. 4-5. Changes in the concentration of ABS (a), COD (b), cell number of m2 and peak area detected at Rt=3.9min of HPLC(d) during the degradation of tin-treated solution of Orange II by co-culture of ABS-degrading consortium and m2.
114
160
30
140 120
25
COD (mg/l)
ABS (mg/l)
35
20 15
(a)
10
80 60 20
0
0
1
(b)
40
5
0
2
3 4 Time (days)
5
6
7
0
1
2
3 4 5 Time (days)
6
7
80000
8.5
70000 8
60000 Peak area
Cell number, log (cfu/ml)
100
7.5 7
40000 30000
(d)
20000
(c)
6.5
50000
10000 0
6 0
1
2
3
4
5
6
0
7
1
2
3
4
5
6
7
Time (days)
Time (days)
Fig. 4-6. Changes in the concentration of ABS(a), COD(b), cell number of m2 (c) and peak area detected at Rt=3.9min of HPLC(d) during the co-culture of ABS-degrading consortium and m2 in the mineral medium containing authentic ABS(50mg/l) and 1A2N (50mg/l).
115
40 35
ABS (mg/l)
30 25 20 15 10
(a)
5 0 0
2
4
6
8
10
Time (days)
70000 60000
Peak area
50000 40000 30000
(b)
20000 10000 0 0
2
4
6
8
10
Time (days)
Fig. 4-7. Changes in the concentration of ABS (a) and peak area detected at Rt=3.9 min of HPLC (b) during the degradation of tin-treated solution of Orange II by the co-culture of ABS-degrading consortium and m2 in the presence of tin.
116
Chapter 5 Decolororization of Orange II by combined use of tin and Pseudomanas aeruginosa m2 in the presence of citrate
5. 1. Introduction In chapter 4, I used tin for the pretreatment of biological degradation of aromatic amines. The decolorization of Orange II by tin was performed under acidic condition, but the neutralization of pH was needed for the subsequent biological treatment [1]. This operation would be difficult in the wastewater process. Nam and Tratnyek reported that zero-valence iron decolorized azo dyes at around neutral pH with a high degree of efficiency [2]. They used HEPES buffer as reaction solvent and oxygen-free condition to avoid the iron process drawbacks such as the coverage of the iron surface by ferric oxide or hydroxide, resulting in reducing the active site [3] and huge sludge production from the hydrolysis of Fe (Ⅲ) ion [4]. The utilization of buffer solution and the establishment of oxygen free condition were not also practical in wastewater treatment. Biological decolorization was considered to be cost effective under proper conditions. Some bacterial strains such as Shewanella decolorationis [5], Pseudomonas luteola [6-8], Pseudomonas sp. [9-11] and several strains of P. aeruginosa were reported as efficient decolorizers of azo dyes [12-17] under static conditions. Glucose as electron donating substrate and yeast extract which includes large amounts of NADH, an electron shuttling coenzyme, were prerequisite for the decolorization [12]. Some bacteria which are capable of decolorization of azo compound synthesized aerobic azoreductase which can reductively cleave the azo group in the presence of oxygen [18]. However, the decolorization of azo dyes using bacterial culture proceeded 117
in most cases under anaerobic or static condition. If the environment is aerobic in shaking or agitating culture, oxygen inhibits the dye reduction mainly because the electron liberated from the oxidation of the electron donors by the cells is preferentially used to reduce oxygen rather than the azo dye [18, 19]. In chapter 4, I performed the chemical decolorization of Orange II followed by biological degradation of produced aromatic amines by an isolated Pseudomonas aeruginosa strain m2 as 1-amino-2-naphthol degrader [1]. In chapters 2 and 3, I described the enhancement of decolorization of azo dyes by PIP tin balls or zero-valence tin in the presence of citric acid [20, 21]. Although the acidic condition (pH
