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Fate of iron oxalates in aqueous solution: The role of temperature, iron species and dissolved oxygen
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JECE 451 1–6 Journal of Environmental Chemical Engineering xxx (2014) xxx–xxx
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Fate of iron oxalates in aqueous solution: The role of temperature, iron species and dissolved oxygen
3 Q1
Gema Pliego * , Juan A. Zazo, José A. Casas, Juan J. Rodríguez
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Chemical Engineering, Faculty of Sciences, Universidad Autónoma de Madrid, Ctra. de Colmenar Km 15, Madrid 28049, Spain
A R T I C L E I N F O
A B S T R A C T
Article history: Received 27 May 2014 Accepted 18 September 2014
The effect of temperature (up to 120 C) and the iron species (Fe2+ or Fe3+) as well as dissolved oxygen on the degradation of oxalic acid in aqueous solutions was studied. The results obtained showed that the fate of oxalate depends on the concentration of ferric ions, which are involved in the formation of oxalate complexes, and the presence of oxygen which oxidizes iron species. By increasing the temperature (120 C), the ferric oxalate complexes undergo a ligand-to-metal electron-transfer reaction, leading to the homolytic cleavage of a FeIII–O coordination bond of the oxalate anion ligand giving rise to ferrous ion, oxalate anion and oxalate radical. The presence of dissolved oxygen enhances the degradation of oxalic acid by oxidizing Fe2+ to Fe3+. This explains the fact that oxalic acid is easily mineralized by Fenton oxidation at high temperature (above 100 C), whereas it proves refractory at lower temperatures. Experiments with a radical scavenger were also carried out to learn about the role of OH radicals. ã 2014 Published by Elsevier Ltd.
Keywords: Oxalic acid Iron oxalates Fenton oxidation High temperature
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1. Introduction
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Oxalic acid is commonly found among the byproducts of the oxidative degradation of a large variety of organic compounds [1]. In particular, it is fairly resistant to mineralization under the operating conditions of the conventional Fenton process [2], one of the most frequently-used advanced oxidation techniques. In spite of its low ecotoxicity it has shown a poor response in biodegradability tests and therefore further abatement with biological treatment appears unlikely [3]. In the Fenton process, oxalic acid can deactivate the iron cations through the formation of highly stable complexes. The formation of iron oxalates and their photo-activity has been widely researched, since the breakdown of oxalates is strongly enhanced in photoassisted AOP [5–7]. Depending on the pH and the concentrations of Fe and oxalate ions, complexes with one, two or three ligands in the coordination sphere of Fe can exist in aqueous solution (Reactions (1)–(6)) [6]. II (1) Fe2þ þ C2 O2 4 $Fe ðC2 OÞ4
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* Corresponding author. Tel.: +34 914975575; fax: +34 914973516. E-mail address: [email protected] (G. Pliego).
2 II FeII ðC2 O4 Þ þ C2 O2 4 $Fe ðC2 O4 Þ2
(2)
þ III Fe3þ þ C2 O2 4 $Fe ðC2 O4 Þ
(3)
III FeIII ðC2 O4 Þþ þ C2 O2 4 $Fe ðC2 O4 Þ2
(4)
2 3 III FeIII ðC2 O4 Þ 2 þ C2 O4 $Fe ðC2 O4 Þ3
(5)
3þ III FeIII ðC2 O4 Þ3 3 þ Fe $Fe2 ðC2 O4 Þ3
(6)
The mineralization of oxalic acid by photocatalytic, ultrasonic, electrochemical and wet air oxidation has been studied [8–12], Q2 although they suffer from important technical and economic drawbacks [13]. In a recent work, we demonstrated that oxalic acid or oxalate could be readily mineralized using Fenton oxidation by increasing the temperature, even after total consumption of H2O2 [14], unlike what occurs at near-room temperature. This study has the double objective of elucidating the fate of ferrous and ferric oxalates in aqueous solution at high temperature (120 C) in the presence and
http://dx.doi.org/10.1016/j.jece.2014.09.013 2213-3437/ ã 2014 Published by Elsevier Ltd.
Please cite this article in press as: G. Pliego, et al., Fate of iron oxalates in aqueous solution: The role of temperature, iron species and dissolved oxygen, J. Environ. Chem. Eng. (2014), http://dx.doi.org/10.1016/j.jece.2014.09.013
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absence of dissolved oxygen and to address the mineralization of oxalic acid by OH radicals in high-temperature Fenton oxidation.
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2. Materials and methods
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2.1. Chemicals
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Oxalic acid (OxAc), +99%; ferrous sulphate (FeSO47H2O), and ferric nitrate ((FeNO3)39H2O), were purchased at Sigma–Aldrich. Hydrogen peroxide (H2O2), analytical grade, w = 30% was supplied by Panreac.
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2.2. Experimental procedure
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The experiments were carried out in a 500 mL stoppered glass reactor (Büchi, inertclave Type I) equipped with a backpressure controller. The reactor was initially loaded with 380 mL of oxalic acid solution (105.3 mg/L, to complete 100 mg/L when all the reactants were added). Once the testing temperature (120 C) was reached, 20 mL of iron aqueous solution were injected. The behaviour of the different iron oxalate complexes was checked by using Fe2+ or Fe3+, either under N2-saturated atmosphere conditions (continuous N2 flow of 50 NmL/min) or in the presence of oxygen (QO2 = 50 NmL/min). Iron doses were varied within the range of 10–100 mg/L. In the Fenton oxidation experiments, 10 mL of the hydrogen peroxide solution and 10 mL of the iron aqueous solution were injected into the reactor once the temperature was equilibrated. The initial concentrations of H2O2 and Fe3+ were 500 and 10 mg/L, respectively. The experiments in presence of a radical scavenger were conducted by adding 5 mL of 20 mM tert-butanol (ButOH) solution. In these experiments, N2 was continuously bubbled into the reactor at 50 NmL/min. The initial pH was around 3 in all cases and was not controlled throughout the experiments. Nevertheless, significant pH alterations were never observed during the experiments. Samples were taken periodically from the reactor and analysed immediately thereafter. Oxidation runs were carried out by triplicate, the standard deviation being less than 5% in all cases. Blank experiments with oxalic acid in the presence and in the absence of O2 or H2O2 were also performed.
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2.3. Analytical methods
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Total organic carbon (TOC) was measured by a Shimadzu model VSCH TOC analyser. Oxalate anions were analysed by ion chromatography with chemical suppression (Metrohm 790 IC) using a conductivity detector. A Metrosep A Supplementary 5-250 column (25 cm length, 4 mm internal diameter) was used as stationary phase and a 3.2 mM Na2CO3 aqueous solution as the mobile phase. The total amount of iron (SFen+) was measured by colourimetric titration using the o-phenantroline method. The concentration of hydrogen peroxide during the Fenton experiments was determined by the colourimetric TiOSO4 method [15].
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3. Results and discussion
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3.1. Fate of iron oxalate complexes at high temperature
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The speciation of dissolved iron in aqueous solution in the presence of oxalic acid depends on the competition between the formation of iron–hydroxo and iron–oxalate complexes [4]. Zuo Q3 and Hoigne [5] reported that under acidic conditions (pH range 3–5), Fe2+ mainly occurs as the hydrated cation [Fe(OH)x]2x, whereas Fe3+–oxalate complexes could be the predominant dissolved species. This latest can be easily confirmed by a spectrophotometer measurement since Fe(III)–oxalate complexes
have absorption bands that extend into the 290–570 nm, being the maximum absorbance at 274 nm [5]. The authors of the above mentioned paper also reported that the amount of oxalic acid oxidized to CO2 was proportional to the amount of Fe2+ in solution, reduced under UV light and inert conditions, according to the following overall Reaction (7): hv
! 2Fe2þ þ ð2n 1ÞC2 O2 2FeðC2 O4 Þð32nÞþ 4 þ 2CO2 n
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(8)
Hence, the degradation of ferric oxalates at high temperature proceeds through a thermal/chemical cycle involving Fe3+, oxalic acid and O2. Increasing the temperature promotes an electrotransfer from a complexing oxalate ligand to the central ferric ion, leading to the homolytic breaking of a FeIII–O coordination bond of the oxalate ligand giving rise to ferrous ion, oxalate anion and oxalate radical anion, according to Reaction (9) [5]. Then, the oxalate radical can reduce another FeIII–oxalate complex according to the Reaction (10). Therefore, the overall mechanism for oxalate mineralization (Reaction (7) can be summarized as the sum of these two Reactions ((9) and (10)). Ta
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(7)
This reaction, though thermodynamically favourable, requires high activation energy. Thus, the reaction can be promoted by photons (hn) or by thermal energy, i.e. increasing the temperature. Fig. 1 shows the effect of FeII and FeIII on the evolution of TOC (corresponding to 100 mg/L of starting oxalic acid) at 120 C under N2 or O2 atmosphere. Previously, those runs were carried out at 50 C and 90 C but the TOC reduction was negligible in all cases (data not shown). The thermal stability of oxalic acid was confirmed in blank experiments at 120 C in the absence of iron. As can be observed, under N2 atmosphere, complexation between Fe2+ and oxalic acid does not occur and the TOC remained unchanged. Contrariwise, with ferric ions up to 55% TOC reduction was observed due to the decomposition of ferric oxalates by Reaction (7). In the presence of O2, the degradation of oxalic acid was significantly enhanced in the case of both iron species. Under these conditions, Fe2+ is oxidized to Fe3+ (Reaction (8)) [15] which would form new complexes, allowing the degradation cycle of ferric oxalate complexes to continue. Fe2 þ O2 ! Fe3þ þ O2
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! Fe2þ þ ðn 1ÞC2 O2 FeðC2 O4 Þð32nÞþ 4 þ C2 O4 n
(9)
Fig. 1. Effect of iron species and dissolved oxygen on the mineralization of oxalic acid ([OxAc]0 = 100 mg/L; [Fen+]0 = 100 mg/L; pH0 3; T0 = 120 C; QN2/O2 = 50 N mL/min).
Please cite this article in press as: G. Pliego, et al., Fate of iron oxalates in aqueous solution: The role of temperature, iron species and dissolved oxygen, J. Environ. Chem. Eng. (2014), http://dx.doi.org/10.1016/j.jece.2014.09.013
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FeðC2 O4 Þð32nÞþ þ C2 O4 ! Fe2þ þ nC2 O2 4 þ 2CO2 n 127 126 128 129 130
Moreover, the oxalate radical released yields CO2 and the radical anion CO2 by Reaction (11). In the presence of oxygen, the oxalate radical also reacts with dissolved O2 to produce the superoxide ion (O2 ) from Reaction (11):
C2 O4 ! CO2 þ CO2
(11)
C2 O4 þ O2 ! 2CO2 þ O2 131 132 133 134 135 136 137 138 139 140 141 142 143 144
When Fe is added to an oxygen-saturated aqueous solution of oxalic acid, two pathways for the mineralization can be proposed. On one hand, Fe2+ is oxidized to Fe3+ (Reaction (8) and the ferric ions tend to form ferric oxalate complexes with the oxalate anions present in the reaction media. On the other hand, to a lesser extent, metal complexation remains unchanged during one-electron transfer between metal complexes and O2, as shown in Reaction (13) [17]. In both cases, mineralization occurs through the decomposition of the ferric oxalate complexes formed in Reactions (9) and (10). This less favoured pathway (Reaction (13)) was confirmed by experiments with pure Fe(II)–oxalate complexes in the presence of O2. Under these conditions, the reaction rate observed for the TOC depletion was lower than the one obtained when iron(II) and oxalic acid was added separately (data not shown). Fe
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(12)
2+
II
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(10)
ðC2 O4 Þð22nÞþ n
III
þ O2 $Fe
ðC2 O4 Þð32nÞþ n
þ
O2
(13)
3.2. Effect of iron concentration and oxygen on the degradation of ferric oxalate complexes The results obtained indicate that the degradation of oxalic acid depends on the presence of Fe3+ in the reaction media. Thus, in order to elucidate the effect of iron concentration on the evolution of TOC under N2 and O2-saturated atmospheres, the Fe3+ dose was varied within the range of 10–100 mg/L, 20 mg/L being the theoretical stoichiometric amount of iron for the formation of FeIII(C2O4)33 with 100 mg/L of oxalic acid. As can be observed in Fig. 2, increasing the iron dose leads to a greater reduction of TOC. Again, the presence of O2 has a significant influence upon TOC evolution. Under N2-saturated atmosphere, Reactions (9) and (10) are the only degradation pathway for ferric–oxalate. Therefore, the
mineralization of oxalic acid was quite slow, becoming higher as the amount of iron increased since the oxalate anions released from Reaction (9) form new complexes with the dissolved iron Fe3 + . The presence of O2 clearly improves TOC reduction as a result of the oxidation of both Fe2+ (Reaction (8)) and oxalate radicals (Reaction (12)).
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3.3. Role of OH radicals in the mineralization of ferric oxalate at high temperature
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Oxalic acid has traditionally been considered as refractory to the Fenton oxidation process. However, at high temperature, it is easily oxidized [14]. To elucidate the role of OH radicals in the oxidation of oxalate ions at high temperature, a set of experiments were carried out with H2O2 in the presence of the radical scavenger tert-butanol (ButOH) under N2-saturated atmosphere. Effective scavenging of OH by ButOH was ensured by working under ButOH excess with respect to oxalate and H2O2, as well as a high [ButOH]: [Fe] ratio. Hence, the OH radicals produced react with ButOH to give the CH2CMe2OH radical. Finally, this species evolves to nonradical products (Reactions (14) and (15)):
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OH þ But OH ! H2 O þ CH2 CMe2 OH
Fe2þ þ H2 O2 ! Fe3þ þ OH þ OH
The OH radicals generated then react with Bu OH by Reactions (14) and (15). In absence of ButOH, OH radicals react with the oxalate ions released from Reaction (9) producing either formic acid or CO2 (Reactions (17) and (18), respectively). Ion chromatography analyses confirmed the presence of formic acid but only in trace amounts, so that the main oxidation pathway of oxalic acid by OH radicals must be direct mineralization.
C2 O2 4 þ OH ! CO2 þ HCOO
Fig. 2. Effect of Fe concentration on oxalic acid mineralization in the presence and in the absence of O2 ([OxAc]0 = 100 mg/L; pH0 3; T0 = 120 C; QN2/O2 = 50 NmL/min).
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(16) t
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(15)
Fig. 3 shows the time-evolution of oxalate concentration and H2O2 conversion at 50 C and 120 C under N2 atmosphere with 10 mg/L Fe3+ in the presence and in the absence of ButOH. For the sake of comparison, the results obtained under O2-saturated atmosphere with the same amount of iron but in absence of H2O2 are also included. As can be seen, the addition of ButOH reduces the oxalate conversion up to values similar to those achieved without H2O2 under O2-saturated atmosphere. This suggests that in the presence of a scavenger, oxalate oxidation occurs only via the decomposition of ferric oxalates according to Reaction (9). Therefore, H2O2 plays the same role as O2 in Reaction (8), oxidizing Fe2+ up to Fe3+ according to the classical Fenton reaction [16]:
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(14)
2 CH2 CMe2 OH ! non radical products
C2 O2 4 þ OH ! CO2 þ CO2 þ OH
3+
3
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(17)
(18)
As can be observed, in the absence of any radical scavenger and under N2 flow, oxalic acid was eliminated as long as there was H2O2 remaining in the reaction media. After that, around 25 mg/L of oxalic acid remained in solution until the end of the experiments. The higher oxalic acid depletion, the higher amount of free iron in solution and therefore the higher H2O2 converted. Moreover, in the absence of ButOH, the OH radicals produced react with oxalic acid as well as with H2O2 by means of the well-known parasitic Fenton reactions ((19) and (20)) increasing the decomposition rate of H2O2. On the contrary, as can be observed in Fig. 3, the presence of the radical scavenger significantly reduces this decomposition rate. Under
Please cite this article in press as: G. Pliego, et al., Fate of iron oxalates in aqueous solution: The role of temperature, iron species and dissolved oxygen, J. Environ. Chem. Eng. (2014), http://dx.doi.org/10.1016/j.jece.2014.09.013
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Fig. 4. Evolution of iron in solution under N2 (A) and O2 (B) flow ([OxAc]0 = 100 mg/L; [H2O2]0 = 500 mg/L; [ButOH]0 = 20 mM; pH0 3; T0 = 120 C; QN2/O2 = 50 N mL/min). Fig. 3. Time-evolution of oxalic acid concentration and H2O2 conversion upon hightemperature Fenton oxidation, with and without a scavenger ([OxAc]0 = 100 mg/L; [Fe3+]0 = 10 mg/L; [H2O2]0 = 500 mg/L; [ButOH]0 = 20 mM; pH0 3; T0 = 120 C; QN2 = 50 N mL/min; experiment without H2O2: QO2 = 50 N ml/min).
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these conditions, OH radicals react with ButOH instead of with oxalic acid which reduces the amount of free iron in solution as well as the H2O2 conversion.
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OH þ H2 O2 ! OOH þ H2 O
OH þ OOH ! H2 O þ O2
(19)
is sufficient oxalic acid in the reaction media. Nevertheless, under N2 flow (Fig. 4A), if the initial iron concentration is well above the stoichiometric ratio, the excess of Fe3+ gives rise to Fe(OH)2+ (Reaction (21)), which is the dominant monomeric FeIII–hydroxo complex at pH between 2.5 and 5 [18,19]. This species may hydrolyse and precipitate as iron hydroxide (Fe(OH)3) [20] according to Reactions (22) and (23). Fe3þ þ H2 O$FeIII ðOHÞ2þ þ Hþ
(21)
FeIII ðOHÞ2þ þ 2H2 O$FeðOHÞ3 þ 2Hþ
(22)
2FeðOHÞ3 $Fe2 O3 #
(23)
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(20)
Temperature plays a key role in the degradation of oxalic acid by Fenton oxidation. Due to the high activation energy of Reaction (9), ferric ions tend to remain in the oxalate complex at 50 C and the generation of OH radicals – and consequently, the oxidation of oxalic acid – is inhibited. In the absence of oxalic acid, the decomposition of H2O2 by means of ferric ions reached up to 90% after 4 h reaction time (data not shown). 3.4. Evolution of iron in solution Fig. 4 depicts the evolution of iron in solution in the experiments carried out under N2 (A) and O2 (B) flow. In both cases, and with sub-stoichiometric iron dose (10 mg/L), the precipitation of iron was negligible since all the ferric ions released from Reaction (9) or produced from the classical Fenton Reaction (16) tend to form new oxalate complexes as long as there
However, as can be observed in Fig. 4A, the concentration of iron in solution increased after a sharp initial drop. According to Reaction (24), the precipitate may be re-dissolved by the oxalates released from Reaction (9) leading to the formation of new iron oxalate complexes [21] and increasing the amount of iron in solution. þ 2H2 O FeðOHÞ3 þ nðC2 O4 Þð32nÞ þ Hþ ! FeIII ðC2 O4 Þð32nÞþ n
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(24)
In the presence of dissolved O2 at iron concentrations above the stoichiometric ratio, the iron precipitated cannot be re-dissolved because the concentration of oxalic acid in the reaction media is
Please cite this article in press as: G. Pliego, et al., Fate of iron oxalates in aqueous solution: The role of temperature, iron species and dissolved oxygen, J. Environ. Chem. Eng. (2014), http://dx.doi.org/10.1016/j.jece.2014.09.013
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Fig. 5. Diagram for the thermal/chemical degradation iron oxalate (black lines: O2 or N2 flow; red lines: O2; purple lines: N2; blue lines: N2 with H2O2; green lines: N2 with Q5 H2O2 in the presence of ButOH. Dotted lines indicate minor pathways). (For interpretation of the references to colour in this figure legend, the reader is referred to the web version of this article.)
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not high enough due to the rapid degradation of oxalate from Reactions (9) and (10) (see Fig. 2). X-ray diffraction of the reddish brown precipitate separated after the experiments revealed the presence of ferric oxide (hematite) without detected impurities.
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3.5. Degradation pathway of iron oxalate at high temperature
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From the results obtained, the diagram in Fig. 5 is proposed for the degradation of iron oxalate at 120 C. At this temperature, an electron transfer from a complexing oxalate ligand to Fe3+ takes place, producing Fe2+, oxalate anion and oxalate anion radical. This last can further reduce a ferric–oxalate ion. In the presence of O2 or H2O2, this Fe2+ is readily oxidized to Fe3+ which would react with oxalate anions to produce new FeIII–oxalate complexes. This ferric complex could be also formed from the oxidation of FeII–oxalate complex by O2. Hydroxyl radicals, produced from the catalytic decomposition of H2O2, oxidize the oxalate anions. Under N2 flow, Fe2+ remains in the reaction media either free or as in hydroxyl or oxalate complexes. In the presence of O2, oxalate radical anions would react with O2 to produce CO2 and O2. If the initial iron concentration is high enough, the excess of iron tends to precipitate as iron oxide, which would be dissolved by oxalate
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anions under inert conditions (N2). All the reactions described can be written for mono-, di- or tri-oxalate ions.
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4. Conclusions
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The effective mineralization of oxalic acid at high temperature in the presence of ferric ions is attributed to the ligand-to-metal electron transfer reaction where the FeIII–oxalate complex yields Fe2 + , oxalate anion andan oxalate anionradical. The experimentalresults show that the rate of mineralization increases with the concentration of ferric ions. Dissolved oxygen plays an important role in both the FeIII–FeII redox cycle and the oxidation of oxalate anion radicals. The oxidation continues as long as there is enough oxalic acid in the reaction media to form new ferric oxalate complexes. Fenton experiments, in the presence and in absence of a radical scavenger, demonstrate that oxalate can be also mineralized by OH radicals. Hydrogen peroxide regenerates Fe3+ from the Fe2+ formed upon the decomposition of the FeIII–oxalate complex. Then, the OH radicals produced from the classical Fenton reaction oxidize the oxalate anions. The solubility of iron at high temperature decreases significantly leading to the precipitation of iron oxides, which are further dissolved by oxalate anions to produce new FeIII–oxalate complexes in the absence of dissolved oxygen (N2 flow).
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Please cite this article in press as: G. Pliego, et al., Fate of iron oxalates in aqueous solution: The role of temperature, iron species and dissolved oxygen, J. Environ. Chem. Eng. (2014), http://dx.doi.org/10.1016/j.jece.2014.09.013
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Acknowledgments
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This research has been supported by the Spanish MICINN through the projects CTQ2008-03988, CTQ2010-14807, and by the CM through the project S2009/AMB-1588.
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References
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[1] R.V. Shende, V.V. Mahajani, Kinetics of wet air oxidation of glyoxalic acid and oxalic acid, Ind. Eng. Chem. Res. 33 (1994) 3125–3130, doi:http://dx.doi.org/ 10.1021/ie00036a030. [2] J.A. Zazo, J.A. Casas, A.F. Mohedano, M.A. Gilarranz, J.J. Rodríguez, Chemical pathway and kinetics of phenol oxidation by Fenton’s reagent, Environ. Sci. Technol. 39 (2005) 9295–9302. 16382955. [3] C.B. Molina, J.A. Zazo, J.A. Casas, J.J. Rodriguez, CWPO of 4-CP and industrial wastewater with Al–Fe pillared clays, Water Sci. Technol.: J. Int. Assoc. Water Pollut. Res. 61 (2010) 2161–2168, doi:http://dx.doi.org/10.2166/wst.2010.151. 20389016. [4] M.E. Balmer, B. Sulzberger, Atrazine degradation in irradiated iron/oxalate systems: effects of pH and oxalate, Environ. Sci. Technol. 33 (1999) 2418–2424, doi:http://dx.doi.org/10.1021/es9808705. [5] Y.G. Zuo, J. Hoigne, Formation of hydrogen peroxide and depletion of oxalic acid in atmospheric water by photolysis of iron(III)–oxalato complexes, Environ. Sci. Technol. 26 (1992) 1014–1022, doi:http://dx.doi.org/10.1021/ es00029a022. [6] B.C. Faust, R.G. Zepp, Photochemistry of aqueous iron(III)–polycarboxylate complexes: roles in the chemistry of atmospheric and surface waters, Environ. Sci. Technol. 27 (1993) 2517–2522, doi:http://dx.doi.org/10.1021/ es00048a032. [7] Y. Zuo, J. Hoigné, Photochemical decomposition of oxalic, glyoxalic and pyruvic acid catalysed by iron in atmospheric waters, Atmos. Environ. 28 (1994) 1231–1239, doi:http://dx.doi.org/10.1016/1352-2310(94)90270-4. [8] M.M. Kosani c, Photocatalytic degradation of oxalic acid over TiO2 power, J. Photochem. Photobiol. A: Chem. 119 (1998) 119–122, doi:http://dx.doi.org/ 10.1016/S1010-6030(98)00407-9. [9] M. Dükkanci, G. Gündüz, Ultrasonic degradation of oxalic acid in aqueous solutions, Ultrason. Sonochem. 13 (2006) 517–522, doi:http://dx.doi.org/ 10.1016/j.ultsonch.2005.10.005. 16352455.
[10] Y.H. Huang, Y.J. Shih, C.H. Liu, Oxalic acid mineralization by electrochemical oxidation processes, J. Hazard. Mater. 188 (2011) 188–192, doi:http://dx.doi. org/10.1016/j.jhazmat.2011.01.091. 21320749. [11] S. Garcia-Segura, E. Brillas, Mineralization of the recalcitrant oxalic and oxamic acids by electrochemical advanced oxidation processes using a boron-doped diamond anode, Water Res. 45 (2011) 2975–2984, doi:http://dx.doi.org/ 10.1016/j.watres.2011.03.017. 21477836. [12] R.P. Rocha, J.P.S. Sousa, A.M.T. Silva, M.F.R. Pereira, J.L. Figueiredo, Catalytic activity and stability of multiwalled carbon nanotubes in catalytic wet air oxidation of oxalic acid: the role of the basic nature induced by the surface chemistry, Appl. Catal. B Environ. 104 (2011) 330–336, doi:http://dx.doi.org/ 10.1016/j.apcatb.2011.03.009. [13] E.M. Rodríguez, B. Núñez, G. Fernández, F.J. Beltrán, Effects of some carboxylic acids on the Fe(III)/UVA photocatalytic oxidation of muconic acid in water, Appl. Catal. B Environ. 89 (2009) 214–222, doi:http://dx.doi.org/10.1016/j. apcatb.2008.11.030. [14] J.A. Zazo, G. Pliego, S. Blasco, J.A. Casas, J.J. Rodriguez, Intensification of the Fenton process by increasing the temperature, Ind. Eng. Chem. Res. 50 (2011) 866–870. [15] G.M. Eisenberg, Colorimetric determination of hydrogen peroxide, Ind. Eng. Chem. Anal. Ed. 15 (1943) 327–328, doi:http://dx.doi.org/10.1021/i560117a011. [16] E.R. Brown, J.D. Mazzarella, Mechanism of oxidation of ferrous polydentate complexes by dioxygen, J. Electroanal. Chem. Interfacial Electrochem. 222 (1987) 173–192, doi:http://dx.doi.org/10.1016/0022-0728(87)80285-1. [17] J.S.B. Park, P.M. Wood, M.J. Davies, B.C. Gilbert, A.C. Whitwood, A kinetic and ESR investigation of iron(II) oxalate oxidation by hydrogen peroxide and dioxygen as a source of hydroxyl radicals, Free Radic. Res. 27 (1997) 447–458, doi:http://dx.doi.org/10.3109/10715769709065785. 9518062. [18] B.C. Faust, J. Hoigné, Photolysis of Fe(III)–hydroxy complexes as sources of OH radicals in clouds, fog and rain, Atmos. Environ. Part A Gen. Top. 24 (1990) 79–89, doi:http://dx.doi.org/10.1016/0960-1686(90)90443-Q. [19] M. Simunovic, H. Kusic, N. Koprivanac, A.L. Bozic, Treatment of simulated industrial wastewater by photo-Fenton process. Part II. The development of mechanistic model, Chem. Eng. J. 173 (2011) 280–289, doi:http://dx.doi.org/ 10.1016/j.cej.2010.09.030. [20] T.D. Waite, Challenges and opportunities in the use of iron in water and wastewater treatment, Rev. Environ. Sci. Bio/Technol. 1 (2002) 9–15, doi: http://dx.doi.org/10.1023/A:1015131528247. [21] U. Schwertmann, Solubility and dissolution of iron oxides, Plant Soil 130 (1991) 1–25, doi:http://dx.doi.org/10.1007/BF00011851.
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Fate of iron oxalates in aqueous solution: The role of temperature, iron species and dissolved oxygen
3 Q1
Gema Pliego * , Juan A. Zazo, José A. Casas, Juan J. Rodríguez
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Chemical Engineering, Faculty of Sciences, Universidad Autónoma de Madrid, Ctra. de Colmenar Km 15, Madrid 28049, Spain
A R T I C L E I N F O
A B S T R A C T
Article history: Received 27 May 2014 Accepted 18 September 2014
The effect of temperature (up to 120 C) and the iron species (Fe2+ or Fe3+) as well as dissolved oxygen on the degradation of oxalic acid in aqueous solutions was studied. The results obtained showed that the fate of oxalate depends on the concentration of ferric ions, which are involved in the formation of oxalate complexes, and the presence of oxygen which oxidizes iron species. By increasing the temperature (120 C), the ferric oxalate complexes undergo a ligand-to-metal electron-transfer reaction, leading to the homolytic cleavage of a FeIII–O coordination bond of the oxalate anion ligand giving rise to ferrous ion, oxalate anion and oxalate radical. The presence of dissolved oxygen enhances the degradation of oxalic acid by oxidizing Fe2+ to Fe3+. This explains the fact that oxalic acid is easily mineralized by Fenton oxidation at high temperature (above 100 C), whereas it proves refractory at lower temperatures. Experiments with a radical scavenger were also carried out to learn about the role of OH radicals. ã 2014 Published by Elsevier Ltd.
Keywords: Oxalic acid Iron oxalates Fenton oxidation High temperature
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1. Introduction
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Oxalic acid is commonly found among the byproducts of the oxidative degradation of a large variety of organic compounds [1]. In particular, it is fairly resistant to mineralization under the operating conditions of the conventional Fenton process [2], one of the most frequently-used advanced oxidation techniques. In spite of its low ecotoxicity it has shown a poor response in biodegradability tests and therefore further abatement with biological treatment appears unlikely [3]. In the Fenton process, oxalic acid can deactivate the iron cations through the formation of highly stable complexes. The formation of iron oxalates and their photo-activity has been widely researched, since the breakdown of oxalates is strongly enhanced in photoassisted AOP [5–7]. Depending on the pH and the concentrations of Fe and oxalate ions, complexes with one, two or three ligands in the coordination sphere of Fe can exist in aqueous solution (Reactions (1)–(6)) [6]. II (1) Fe2þ þ C2 O2 4 $Fe ðC2 OÞ4
7 8 9 10 11 12 13 14 15 16 17 18 19 20 21
* Corresponding author. Tel.: +34 914975575; fax: +34 914973516. E-mail address: [email protected] (G. Pliego).
2 II FeII ðC2 O4 Þ þ C2 O2 4 $Fe ðC2 O4 Þ2
(2)
þ III Fe3þ þ C2 O2 4 $Fe ðC2 O4 Þ
(3)
III FeIII ðC2 O4 Þþ þ C2 O2 4 $Fe ðC2 O4 Þ2
(4)
2 3 III FeIII ðC2 O4 Þ 2 þ C2 O4 $Fe ðC2 O4 Þ3
(5)
3þ III FeIII ðC2 O4 Þ3 3 þ Fe $Fe2 ðC2 O4 Þ3
(6)
The mineralization of oxalic acid by photocatalytic, ultrasonic, electrochemical and wet air oxidation has been studied [8–12], Q2 although they suffer from important technical and economic drawbacks [13]. In a recent work, we demonstrated that oxalic acid or oxalate could be readily mineralized using Fenton oxidation by increasing the temperature, even after total consumption of H2O2 [14], unlike what occurs at near-room temperature. This study has the double objective of elucidating the fate of ferrous and ferric oxalates in aqueous solution at high temperature (120 C) in the presence and
http://dx.doi.org/10.1016/j.jece.2014.09.013 2213-3437/ ã 2014 Published by Elsevier Ltd.
Please cite this article in press as: G. Pliego, et al., Fate of iron oxalates in aqueous solution: The role of temperature, iron species and dissolved oxygen, J. Environ. Chem. Eng. (2014), http://dx.doi.org/10.1016/j.jece.2014.09.013
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absence of dissolved oxygen and to address the mineralization of oxalic acid by OH radicals in high-temperature Fenton oxidation.
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2. Materials and methods
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2.1. Chemicals
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Oxalic acid (OxAc), +99%; ferrous sulphate (FeSO47H2O), and ferric nitrate ((FeNO3)39H2O), were purchased at Sigma–Aldrich. Hydrogen peroxide (H2O2), analytical grade, w = 30% was supplied by Panreac.
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2.2. Experimental procedure
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The experiments were carried out in a 500 mL stoppered glass reactor (Büchi, inertclave Type I) equipped with a backpressure controller. The reactor was initially loaded with 380 mL of oxalic acid solution (105.3 mg/L, to complete 100 mg/L when all the reactants were added). Once the testing temperature (120 C) was reached, 20 mL of iron aqueous solution were injected. The behaviour of the different iron oxalate complexes was checked by using Fe2+ or Fe3+, either under N2-saturated atmosphere conditions (continuous N2 flow of 50 NmL/min) or in the presence of oxygen (QO2 = 50 NmL/min). Iron doses were varied within the range of 10–100 mg/L. In the Fenton oxidation experiments, 10 mL of the hydrogen peroxide solution and 10 mL of the iron aqueous solution were injected into the reactor once the temperature was equilibrated. The initial concentrations of H2O2 and Fe3+ were 500 and 10 mg/L, respectively. The experiments in presence of a radical scavenger were conducted by adding 5 mL of 20 mM tert-butanol (ButOH) solution. In these experiments, N2 was continuously bubbled into the reactor at 50 NmL/min. The initial pH was around 3 in all cases and was not controlled throughout the experiments. Nevertheless, significant pH alterations were never observed during the experiments. Samples were taken periodically from the reactor and analysed immediately thereafter. Oxidation runs were carried out by triplicate, the standard deviation being less than 5% in all cases. Blank experiments with oxalic acid in the presence and in the absence of O2 or H2O2 were also performed.
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2.3. Analytical methods
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Total organic carbon (TOC) was measured by a Shimadzu model VSCH TOC analyser. Oxalate anions were analysed by ion chromatography with chemical suppression (Metrohm 790 IC) using a conductivity detector. A Metrosep A Supplementary 5-250 column (25 cm length, 4 mm internal diameter) was used as stationary phase and a 3.2 mM Na2CO3 aqueous solution as the mobile phase. The total amount of iron (SFen+) was measured by colourimetric titration using the o-phenantroline method. The concentration of hydrogen peroxide during the Fenton experiments was determined by the colourimetric TiOSO4 method [15].
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3. Results and discussion
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3.1. Fate of iron oxalate complexes at high temperature
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The speciation of dissolved iron in aqueous solution in the presence of oxalic acid depends on the competition between the formation of iron–hydroxo and iron–oxalate complexes [4]. Zuo Q3 and Hoigne [5] reported that under acidic conditions (pH range 3–5), Fe2+ mainly occurs as the hydrated cation [Fe(OH)x]2x, whereas Fe3+–oxalate complexes could be the predominant dissolved species. This latest can be easily confirmed by a spectrophotometer measurement since Fe(III)–oxalate complexes
have absorption bands that extend into the 290–570 nm, being the maximum absorbance at 274 nm [5]. The authors of the above mentioned paper also reported that the amount of oxalic acid oxidized to CO2 was proportional to the amount of Fe2+ in solution, reduced under UV light and inert conditions, according to the following overall Reaction (7): hv
! 2Fe2þ þ ð2n 1ÞC2 O2 2FeðC2 O4 Þð32nÞþ 4 þ 2CO2 n
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96 95 97 98 99 100 101 102 103 104 105 106 107 108 109 110 111 112 113
(8)
Hence, the degradation of ferric oxalates at high temperature proceeds through a thermal/chemical cycle involving Fe3+, oxalic acid and O2. Increasing the temperature promotes an electrotransfer from a complexing oxalate ligand to the central ferric ion, leading to the homolytic breaking of a FeIII–O coordination bond of the oxalate ligand giving rise to ferrous ion, oxalate anion and oxalate radical anion, according to Reaction (9) [5]. Then, the oxalate radical can reduce another FeIII–oxalate complex according to the Reaction (10). Therefore, the overall mechanism for oxalate mineralization (Reaction (7) can be summarized as the sum of these two Reactions ((9) and (10)). Ta
90
(7)
This reaction, though thermodynamically favourable, requires high activation energy. Thus, the reaction can be promoted by photons (hn) or by thermal energy, i.e. increasing the temperature. Fig. 1 shows the effect of FeII and FeIII on the evolution of TOC (corresponding to 100 mg/L of starting oxalic acid) at 120 C under N2 or O2 atmosphere. Previously, those runs were carried out at 50 C and 90 C but the TOC reduction was negligible in all cases (data not shown). The thermal stability of oxalic acid was confirmed in blank experiments at 120 C in the absence of iron. As can be observed, under N2 atmosphere, complexation between Fe2+ and oxalic acid does not occur and the TOC remained unchanged. Contrariwise, with ferric ions up to 55% TOC reduction was observed due to the decomposition of ferric oxalates by Reaction (7). In the presence of O2, the degradation of oxalic acid was significantly enhanced in the case of both iron species. Under these conditions, Fe2+ is oxidized to Fe3+ (Reaction (8)) [15] which would form new complexes, allowing the degradation cycle of ferric oxalate complexes to continue. Fe2 þ O2 ! Fe3þ þ O2
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! Fe2þ þ ðn 1ÞC2 O2 FeðC2 O4 Þð32nÞþ 4 þ C2 O4 n
(9)
Fig. 1. Effect of iron species and dissolved oxygen on the mineralization of oxalic acid ([OxAc]0 = 100 mg/L; [Fen+]0 = 100 mg/L; pH0 3; T0 = 120 C; QN2/O2 = 50 N mL/min).
Please cite this article in press as: G. Pliego, et al., Fate of iron oxalates in aqueous solution: The role of temperature, iron species and dissolved oxygen, J. Environ. Chem. Eng. (2014), http://dx.doi.org/10.1016/j.jece.2014.09.013
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FeðC2 O4 Þð32nÞþ þ C2 O4 ! Fe2þ þ nC2 O2 4 þ 2CO2 n 127 126 128 129 130
Moreover, the oxalate radical released yields CO2 and the radical anion CO2 by Reaction (11). In the presence of oxygen, the oxalate radical also reacts with dissolved O2 to produce the superoxide ion (O2 ) from Reaction (11):
C2 O4 ! CO2 þ CO2
(11)
C2 O4 þ O2 ! 2CO2 þ O2 131 132 133 134 135 136 137 138 139 140 141 142 143 144
When Fe is added to an oxygen-saturated aqueous solution of oxalic acid, two pathways for the mineralization can be proposed. On one hand, Fe2+ is oxidized to Fe3+ (Reaction (8) and the ferric ions tend to form ferric oxalate complexes with the oxalate anions present in the reaction media. On the other hand, to a lesser extent, metal complexation remains unchanged during one-electron transfer between metal complexes and O2, as shown in Reaction (13) [17]. In both cases, mineralization occurs through the decomposition of the ferric oxalate complexes formed in Reactions (9) and (10). This less favoured pathway (Reaction (13)) was confirmed by experiments with pure Fe(II)–oxalate complexes in the presence of O2. Under these conditions, the reaction rate observed for the TOC depletion was lower than the one obtained when iron(II) and oxalic acid was added separately (data not shown). Fe
146 147 148 149 150 151 152 153 154 155 156 157
(12)
2+
II
145
(10)
ðC2 O4 Þð22nÞþ n
III
þ O2 $Fe
ðC2 O4 Þð32nÞþ n
þ
O2
(13)
3.2. Effect of iron concentration and oxygen on the degradation of ferric oxalate complexes The results obtained indicate that the degradation of oxalic acid depends on the presence of Fe3+ in the reaction media. Thus, in order to elucidate the effect of iron concentration on the evolution of TOC under N2 and O2-saturated atmospheres, the Fe3+ dose was varied within the range of 10–100 mg/L, 20 mg/L being the theoretical stoichiometric amount of iron for the formation of FeIII(C2O4)33 with 100 mg/L of oxalic acid. As can be observed in Fig. 2, increasing the iron dose leads to a greater reduction of TOC. Again, the presence of O2 has a significant influence upon TOC evolution. Under N2-saturated atmosphere, Reactions (9) and (10) are the only degradation pathway for ferric–oxalate. Therefore, the
mineralization of oxalic acid was quite slow, becoming higher as the amount of iron increased since the oxalate anions released from Reaction (9) form new complexes with the dissolved iron Fe3 + . The presence of O2 clearly improves TOC reduction as a result of the oxidation of both Fe2+ (Reaction (8)) and oxalate radicals (Reaction (12)).
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3.3. Role of OH radicals in the mineralization of ferric oxalate at high temperature
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Oxalic acid has traditionally been considered as refractory to the Fenton oxidation process. However, at high temperature, it is easily oxidized [14]. To elucidate the role of OH radicals in the oxidation of oxalate ions at high temperature, a set of experiments were carried out with H2O2 in the presence of the radical scavenger tert-butanol (ButOH) under N2-saturated atmosphere. Effective scavenging of OH by ButOH was ensured by working under ButOH excess with respect to oxalate and H2O2, as well as a high [ButOH]: [Fe] ratio. Hence, the OH radicals produced react with ButOH to give the CH2CMe2OH radical. Finally, this species evolves to nonradical products (Reactions (14) and (15)):
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OH þ But OH ! H2 O þ CH2 CMe2 OH
Fe2þ þ H2 O2 ! Fe3þ þ OH þ OH
The OH radicals generated then react with Bu OH by Reactions (14) and (15). In absence of ButOH, OH radicals react with the oxalate ions released from Reaction (9) producing either formic acid or CO2 (Reactions (17) and (18), respectively). Ion chromatography analyses confirmed the presence of formic acid but only in trace amounts, so that the main oxidation pathway of oxalic acid by OH radicals must be direct mineralization.
C2 O2 4 þ OH ! CO2 þ HCOO
Fig. 2. Effect of Fe concentration on oxalic acid mineralization in the presence and in the absence of O2 ([OxAc]0 = 100 mg/L; pH0 3; T0 = 120 C; QN2/O2 = 50 NmL/min).
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(16) t
160
(15)
Fig. 3 shows the time-evolution of oxalate concentration and H2O2 conversion at 50 C and 120 C under N2 atmosphere with 10 mg/L Fe3+ in the presence and in the absence of ButOH. For the sake of comparison, the results obtained under O2-saturated atmosphere with the same amount of iron but in absence of H2O2 are also included. As can be seen, the addition of ButOH reduces the oxalate conversion up to values similar to those achieved without H2O2 under O2-saturated atmosphere. This suggests that in the presence of a scavenger, oxalate oxidation occurs only via the decomposition of ferric oxalates according to Reaction (9). Therefore, H2O2 plays the same role as O2 in Reaction (8), oxidizing Fe2+ up to Fe3+ according to the classical Fenton reaction [16]:
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(14)
2 CH2 CMe2 OH ! non radical products
C2 O2 4 þ OH ! CO2 þ CO2 þ OH
3+
3
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(17)
(18)
As can be observed, in the absence of any radical scavenger and under N2 flow, oxalic acid was eliminated as long as there was H2O2 remaining in the reaction media. After that, around 25 mg/L of oxalic acid remained in solution until the end of the experiments. The higher oxalic acid depletion, the higher amount of free iron in solution and therefore the higher H2O2 converted. Moreover, in the absence of ButOH, the OH radicals produced react with oxalic acid as well as with H2O2 by means of the well-known parasitic Fenton reactions ((19) and (20)) increasing the decomposition rate of H2O2. On the contrary, as can be observed in Fig. 3, the presence of the radical scavenger significantly reduces this decomposition rate. Under
Please cite this article in press as: G. Pliego, et al., Fate of iron oxalates in aqueous solution: The role of temperature, iron species and dissolved oxygen, J. Environ. Chem. Eng. (2014), http://dx.doi.org/10.1016/j.jece.2014.09.013
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Fig. 4. Evolution of iron in solution under N2 (A) and O2 (B) flow ([OxAc]0 = 100 mg/L; [H2O2]0 = 500 mg/L; [ButOH]0 = 20 mM; pH0 3; T0 = 120 C; QN2/O2 = 50 N mL/min). Fig. 3. Time-evolution of oxalic acid concentration and H2O2 conversion upon hightemperature Fenton oxidation, with and without a scavenger ([OxAc]0 = 100 mg/L; [Fe3+]0 = 10 mg/L; [H2O2]0 = 500 mg/L; [ButOH]0 = 20 mM; pH0 3; T0 = 120 C; QN2 = 50 N mL/min; experiment without H2O2: QO2 = 50 N ml/min).
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these conditions, OH radicals react with ButOH instead of with oxalic acid which reduces the amount of free iron in solution as well as the H2O2 conversion.
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OH þ H2 O2 ! OOH þ H2 O
OH þ OOH ! H2 O þ O2
(19)
is sufficient oxalic acid in the reaction media. Nevertheless, under N2 flow (Fig. 4A), if the initial iron concentration is well above the stoichiometric ratio, the excess of Fe3+ gives rise to Fe(OH)2+ (Reaction (21)), which is the dominant monomeric FeIII–hydroxo complex at pH between 2.5 and 5 [18,19]. This species may hydrolyse and precipitate as iron hydroxide (Fe(OH)3) [20] according to Reactions (22) and (23). Fe3þ þ H2 O$FeIII ðOHÞ2þ þ Hþ
(21)
FeIII ðOHÞ2þ þ 2H2 O$FeðOHÞ3 þ 2Hþ
(22)
2FeðOHÞ3 $Fe2 O3 #
(23)
227 228 229 230 231 232 233
(20)
Temperature plays a key role in the degradation of oxalic acid by Fenton oxidation. Due to the high activation energy of Reaction (9), ferric ions tend to remain in the oxalate complex at 50 C and the generation of OH radicals – and consequently, the oxidation of oxalic acid – is inhibited. In the absence of oxalic acid, the decomposition of H2O2 by means of ferric ions reached up to 90% after 4 h reaction time (data not shown). 3.4. Evolution of iron in solution Fig. 4 depicts the evolution of iron in solution in the experiments carried out under N2 (A) and O2 (B) flow. In both cases, and with sub-stoichiometric iron dose (10 mg/L), the precipitation of iron was negligible since all the ferric ions released from Reaction (9) or produced from the classical Fenton Reaction (16) tend to form new oxalate complexes as long as there
However, as can be observed in Fig. 4A, the concentration of iron in solution increased after a sharp initial drop. According to Reaction (24), the precipitate may be re-dissolved by the oxalates released from Reaction (9) leading to the formation of new iron oxalate complexes [21] and increasing the amount of iron in solution. þ 2H2 O FeðOHÞ3 þ nðC2 O4 Þð32nÞ þ Hþ ! FeIII ðC2 O4 Þð32nÞþ n
235 234 236 237 238 239 240
(24)
In the presence of dissolved O2 at iron concentrations above the stoichiometric ratio, the iron precipitated cannot be re-dissolved because the concentration of oxalic acid in the reaction media is
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Fig. 5. Diagram for the thermal/chemical degradation iron oxalate (black lines: O2 or N2 flow; red lines: O2; purple lines: N2; blue lines: N2 with H2O2; green lines: N2 with Q5 H2O2 in the presence of ButOH. Dotted lines indicate minor pathways). (For interpretation of the references to colour in this figure legend, the reader is referred to the web version of this article.)
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not high enough due to the rapid degradation of oxalate from Reactions (9) and (10) (see Fig. 2). X-ray diffraction of the reddish brown precipitate separated after the experiments revealed the presence of ferric oxide (hematite) without detected impurities.
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3.5. Degradation pathway of iron oxalate at high temperature
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From the results obtained, the diagram in Fig. 5 is proposed for the degradation of iron oxalate at 120 C. At this temperature, an electron transfer from a complexing oxalate ligand to Fe3+ takes place, producing Fe2+, oxalate anion and oxalate anion radical. This last can further reduce a ferric–oxalate ion. In the presence of O2 or H2O2, this Fe2+ is readily oxidized to Fe3+ which would react with oxalate anions to produce new FeIII–oxalate complexes. This ferric complex could be also formed from the oxidation of FeII–oxalate complex by O2. Hydroxyl radicals, produced from the catalytic decomposition of H2O2, oxidize the oxalate anions. Under N2 flow, Fe2+ remains in the reaction media either free or as in hydroxyl or oxalate complexes. In the presence of O2, oxalate radical anions would react with O2 to produce CO2 and O2. If the initial iron concentration is high enough, the excess of iron tends to precipitate as iron oxide, which would be dissolved by oxalate
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anions under inert conditions (N2). All the reactions described can be written for mono-, di- or tri-oxalate ions.
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4. Conclusions
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The effective mineralization of oxalic acid at high temperature in the presence of ferric ions is attributed to the ligand-to-metal electron transfer reaction where the FeIII–oxalate complex yields Fe2 + , oxalate anion andan oxalate anionradical. The experimentalresults show that the rate of mineralization increases with the concentration of ferric ions. Dissolved oxygen plays an important role in both the FeIII–FeII redox cycle and the oxidation of oxalate anion radicals. The oxidation continues as long as there is enough oxalic acid in the reaction media to form new ferric oxalate complexes. Fenton experiments, in the presence and in absence of a radical scavenger, demonstrate that oxalate can be also mineralized by OH radicals. Hydrogen peroxide regenerates Fe3+ from the Fe2+ formed upon the decomposition of the FeIII–oxalate complex. Then, the OH radicals produced from the classical Fenton reaction oxidize the oxalate anions. The solubility of iron at high temperature decreases significantly leading to the precipitation of iron oxides, which are further dissolved by oxalate anions to produce new FeIII–oxalate complexes in the absence of dissolved oxygen (N2 flow).
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Please cite this article in press as: G. Pliego, et al., Fate of iron oxalates in aqueous solution: The role of temperature, iron species and dissolved oxygen, J. Environ. Chem. Eng. (2014), http://dx.doi.org/10.1016/j.jece.2014.09.013
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Acknowledgments
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This research has been supported by the Spanish MICINN through the projects CTQ2008-03988, CTQ2010-14807, and by the CM through the project S2009/AMB-1588.
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References
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[1] R.V. Shende, V.V. Mahajani, Kinetics of wet air oxidation of glyoxalic acid and oxalic acid, Ind. Eng. Chem. Res. 33 (1994) 3125–3130, doi:http://dx.doi.org/ 10.1021/ie00036a030. [2] J.A. Zazo, J.A. Casas, A.F. Mohedano, M.A. Gilarranz, J.J. Rodríguez, Chemical pathway and kinetics of phenol oxidation by Fenton’s reagent, Environ. Sci. Technol. 39 (2005) 9295–9302. 16382955. [3] C.B. Molina, J.A. Zazo, J.A. Casas, J.J. Rodriguez, CWPO of 4-CP and industrial wastewater with Al–Fe pillared clays, Water Sci. Technol.: J. Int. Assoc. Water Pollut. Res. 61 (2010) 2161–2168, doi:http://dx.doi.org/10.2166/wst.2010.151. 20389016. [4] M.E. Balmer, B. Sulzberger, Atrazine degradation in irradiated iron/oxalate systems: effects of pH and oxalate, Environ. Sci. Technol. 33 (1999) 2418–2424, doi:http://dx.doi.org/10.1021/es9808705. [5] Y.G. Zuo, J. Hoigne, Formation of hydrogen peroxide and depletion of oxalic acid in atmospheric water by photolysis of iron(III)–oxalato complexes, Environ. Sci. Technol. 26 (1992) 1014–1022, doi:http://dx.doi.org/10.1021/ es00029a022. [6] B.C. Faust, R.G. Zepp, Photochemistry of aqueous iron(III)–polycarboxylate complexes: roles in the chemistry of atmospheric and surface waters, Environ. Sci. Technol. 27 (1993) 2517–2522, doi:http://dx.doi.org/10.1021/ es00048a032. [7] Y. Zuo, J. Hoigné, Photochemical decomposition of oxalic, glyoxalic and pyruvic acid catalysed by iron in atmospheric waters, Atmos. Environ. 28 (1994) 1231–1239, doi:http://dx.doi.org/10.1016/1352-2310(94)90270-4. [8] M.M. Kosani c, Photocatalytic degradation of oxalic acid over TiO2 power, J. Photochem. Photobiol. A: Chem. 119 (1998) 119–122, doi:http://dx.doi.org/ 10.1016/S1010-6030(98)00407-9. [9] M. Dükkanci, G. Gündüz, Ultrasonic degradation of oxalic acid in aqueous solutions, Ultrason. Sonochem. 13 (2006) 517–522, doi:http://dx.doi.org/ 10.1016/j.ultsonch.2005.10.005. 16352455.
[10] Y.H. Huang, Y.J. Shih, C.H. Liu, Oxalic acid mineralization by electrochemical oxidation processes, J. Hazard. Mater. 188 (2011) 188–192, doi:http://dx.doi. org/10.1016/j.jhazmat.2011.01.091. 21320749. [11] S. Garcia-Segura, E. Brillas, Mineralization of the recalcitrant oxalic and oxamic acids by electrochemical advanced oxidation processes using a boron-doped diamond anode, Water Res. 45 (2011) 2975–2984, doi:http://dx.doi.org/ 10.1016/j.watres.2011.03.017. 21477836. [12] R.P. Rocha, J.P.S. Sousa, A.M.T. Silva, M.F.R. Pereira, J.L. Figueiredo, Catalytic activity and stability of multiwalled carbon nanotubes in catalytic wet air oxidation of oxalic acid: the role of the basic nature induced by the surface chemistry, Appl. Catal. B Environ. 104 (2011) 330–336, doi:http://dx.doi.org/ 10.1016/j.apcatb.2011.03.009. [13] E.M. Rodríguez, B. Núñez, G. Fernández, F.J. Beltrán, Effects of some carboxylic acids on the Fe(III)/UVA photocatalytic oxidation of muconic acid in water, Appl. Catal. B Environ. 89 (2009) 214–222, doi:http://dx.doi.org/10.1016/j. apcatb.2008.11.030. [14] J.A. Zazo, G. Pliego, S. Blasco, J.A. Casas, J.J. Rodriguez, Intensification of the Fenton process by increasing the temperature, Ind. Eng. Chem. Res. 50 (2011) 866–870. [15] G.M. Eisenberg, Colorimetric determination of hydrogen peroxide, Ind. Eng. Chem. Anal. Ed. 15 (1943) 327–328, doi:http://dx.doi.org/10.1021/i560117a011. [16] E.R. Brown, J.D. Mazzarella, Mechanism of oxidation of ferrous polydentate complexes by dioxygen, J. Electroanal. Chem. Interfacial Electrochem. 222 (1987) 173–192, doi:http://dx.doi.org/10.1016/0022-0728(87)80285-1. [17] J.S.B. Park, P.M. Wood, M.J. Davies, B.C. Gilbert, A.C. Whitwood, A kinetic and ESR investigation of iron(II) oxalate oxidation by hydrogen peroxide and dioxygen as a source of hydroxyl radicals, Free Radic. Res. 27 (1997) 447–458, doi:http://dx.doi.org/10.3109/10715769709065785. 9518062. [18] B.C. Faust, J. Hoigné, Photolysis of Fe(III)–hydroxy complexes as sources of OH radicals in clouds, fog and rain, Atmos. Environ. Part A Gen. Top. 24 (1990) 79–89, doi:http://dx.doi.org/10.1016/0960-1686(90)90443-Q. [19] M. Simunovic, H. Kusic, N. Koprivanac, A.L. Bozic, Treatment of simulated industrial wastewater by photo-Fenton process. Part II. The development of mechanistic model, Chem. Eng. J. 173 (2011) 280–289, doi:http://dx.doi.org/ 10.1016/j.cej.2010.09.030. [20] T.D. Waite, Challenges and opportunities in the use of iron in water and wastewater treatment, Rev. Environ. Sci. Bio/Technol. 1 (2002) 9–15, doi: http://dx.doi.org/10.1023/A:1015131528247. [21] U. Schwertmann, Solubility and dissolution of iron oxides, Plant Soil 130 (1991) 1–25, doi:http://dx.doi.org/10.1007/BF00011851.
Please cite this article in press as: G. Pliego, et al., Fate of iron oxalates in aqueous solution: The role of temperature, iron species and dissolved oxygen, J. Environ. Chem. Eng. (2014), http://dx.doi.org/10.1016/j.jece.2014.09.013
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