Article pubs.acs.org/IECR
Applied Kinetics Aspects of Ferric EDTA Complex Reduction with Metal Powder Tomasz T. Suchecki,†,⊥ Barbara Mathews,‡,⊥ Adam W. Augustyniak,*,† and Hidehiro Kumazawa§,⊥ †
Faculty of Environmental Engineering, Wroclaw University of Technology, Wybrzeze Wyspianskiego 27, 50-379 Wroclaw, Poland Institute of Environmental Engineering, Polish Academy of Sciences, M. Curie-Sklodowskiej 34, 41-819 Zabrze, Poland § University of Toyama, Toyama, Japan ‡
ABSTRACT: Nitrogen and sulfur oxides (NOx and SOx) emitted from fossil-fuel-fired facilities can be simultaneously absorbed into aqueous mixed solutions of sulfite and FeIIL, a ferrous ion coordinated to a ligand L (a chelate compound). Usually, the ligand L represents an anion of ethylenediaminetetraacetic acid (EDTA or edta). Nitrogen oxide (NO) sparingly soluble in water is promptly bound to the [Fe II(edta)(H 2O)]2− complex to form the [FeII(edta)(NO] 2−compound. However, [FeII(edta)(H2O]2− chelate is oxidized to form [FeIII(edta)(H2O)]− by NO, NO2, and O2, normally coexisting in flue gases. Because the ferric complex does not react with NO, the suppression of these undesired oxidation processes is a very important technological problem to be overcome. Thus, in this work, the reduction kinetics of ferric ion by metal was discussed on the basis of the kinetic data regarding the ferric ion reduction in aqueous solutions of [FeIII(edta)(H2O)]− containing such pulverized metals as aluminum, tin, and zinc. The rate expression for the reduction of [FeIII(edta)(H2O]− by a metal powder was derived in terms of the rate-determining step approximation method. Laboratory-scale experiments were conducted in a recirculation-type glass reactor filled with 250 cm3 of solution. The influence of the following parameters on the reduction rate was investigated: (1) [FeIII(edta)(H2O)]0− = 0.100−0.010 M, (2) [metal]0/[FeIII(edta)(H2O)]0− = 10−100, (3) pH0 = 3−9, and (4) temperature = 20−80 °C. Zinc and aluminum powders were found to be effective reducing agents for ferric ions coordinated to EDTA. Similar experiments were carried out using a working solution in a simultaneous NO and SO2 removal process as a solvent of [FeIII(edta)(H2O)]− solution.
1. INTRODUCTION All commercialized air protection technologies for respective removal of NO and SO2 emitted from stationary combustion facilities (i.e., selective catalytic reduction for NO removal and limestone aqueous slurry for SO2 removal) seem to be expensive from the viewpoints of both investment and operating costs. Therefore, there has still been pressure from the industry to develop new cost-effective methods to improve existing technologies. Because the scrubbing solution for wet NO removal can also remove SO2, the wet scrubbing process is one of the potential candidates for simultaneous NO and SO2. For example, by use of liquid-phase reaction sequences in which both dissolved NO and SO2 participate, the simultaneous removal of both gases can effectively be achieved in a single step or equipment.1 Thus, simultaneous removal processes of NO and SO2 as an alternative to separate removals have been of interest for many research groups worldwide working in the air pollution control field. One of them is the chelate method, whose theoretical fundamentals have been investigated since the 1970s by several research laboratories. However, even now, there has been continuous interested in this process.2−5 The chelate method is based on the combined absorption of SO2 and NOx into aqueous solutions of sulfite and FeIIL, where the ferrous ion is coordinated to a ligand (a chelate compound). Ethylenediaminetetraacetic acid (EDTA or edta) and other compounds were investigated for this purpose as this ligand. Kinetic data and mechanisms of SO2 and NO absorption in these systems have been reported elsewhere.6−9 Briefly, only [FeII(edta)(H2O)]2− (ferrous EDTA chelate) has © 2014 American Chemical Society
an ability to bind NO (sparingly soluble in water) in the liquid phase, contrary to [FeIII(edta)(H2O)]− (ferric EDTA chelate). The Fe2+ ions coordinated to EDTA are continuously oxidized by NO, NO2, and O2 in the flue gas during the SO2/NO absorption process.10−15 Finally, the absorption solution tends to become completely inactive in relation to NO. Therefore, to maintain the degree of NO removal at a high level, it is necessary to keep the concentration of FeII in the absorption solution as high as possible. Practically some portion of the absorption solution should be taken out of the absorber and regenerated by reduction with a coexisting reducing agent. Thus, the regeneration of the absorption solution, that is, the reduction of [FeIII(edta)(H2O]− back to [FeII(edta)(H2O)]2−, is a very important part of the method. Some agents such as the dithionite ion (S2O42−),16 hydrazine (N2H4),17 polyphenolic compound,18 and a chemical absorption−biological reduction (BioDENOx)19−25 were found to be efficient reducing agents of the ferric ion in [FeIII(edta)(H2O]− at ambient temperature. Activated carbon26−28 also acts as an effective reduction agent in this system. At present, only the sulfite ion29−32 has been widely examined as a reducing agent to the ferrous ion in this system. There has been no idea to employ gaseous hydrogen sulfite (H2S),33 a very efficient reducing agent to [FeIII(edta)(H2O]− in the systems of desulfurization of natural gases, for Received: Revised: Accepted: Published: 14234
May 22, 2014 August 10, 2014 August 22, 2014 August 22, 2014 dx.doi.org/10.1021/ie502100h | Ind. Eng. Chem. Res. 2014, 53, 14234−14240
Industrial & Engineering Chemistry Research
Article
d[Fe 2 +] = k6[KM] dt
this purpose. In this work, the rate expression for [FeIII(edta)(H2O)]− reduction by metal was derived by use of rate-determining step approximation. Next, kinetic investigation of reactions of [FeIII(edta)(H2O)]− with aluminum, tin, and zinc powders were presented under the following conditions: (1) [FeIII(edta)(H2O)]0− = 0.100−0.010 M; (2) [metal]0/[FeIII(edta)(H2O)]0− = 10−100; (3) pH0 = 3−9; (4) temperature = 20−80 °C. To represent the significance of the present research, similar experiments were carried out using a real working solution in a simultaneous NO and SO2 removal process as a solvent of [FeIII(edta)(H2O)]− solution. The working solution was found to compare to or rather be superior to the redistilled water.
Elimination of eqs 9, 7, and 4 leads to d[Fe 2 +] [Fe IIIL][M] = k6K 2K3K4 dt ([Fe 2 +][OH−][L])
By substitution of [Fe ] from eq 11 into eq 4, one gets
Fe3 + + L = Fe IIIL
d[Fe 2 +] [Fe IIIL][M] = k6K 2K3K4K5 dt ([Fe IIL][OH−])
(15)
d[Fe 2 +] [Fe IIIL] = ka dt Fe IIL
(16)
or
where ka = k6K 2K3K4K5
(1)
where L stands for ligand (EDTA). For divalent and trivalent metals, m is equal to 1/2 and 1/3, respectively. The reaction rate expression can be derived by assuming the following elementary steps involving a single rate-determining step like in our previous papers16 K1
(14)
2+
2. CHEMISTRY AND KINETIC OF REDUCTION OF FEIII BY METAL POWDER The reduction of ferric ion (FeIII) by metal powder (M) with coexisting EDTA can be can be written stoichiometrically as Fe IIIL + M → Fe IIL + M(1/ m) +
(13)
[M] [OH−]
(17)
ka refers to the apparent reaction rate constant [mol/(L s)] and increases with decreasing pH. Integration of eq 16 with initial conditions at t = 0, [FeIIIL] = a0, and [FeII] = 0 yields ⎧ a0 ⎫ ⎬ − x = kat a0 ln⎨ ⎩ a0 − x ⎭
(2)
(18) II
where x refers to the concentration of Fe L, which can approximately be equal to the concentration of total converted FeII, at any reaction time. Accordingly, the term a0 − x denotes the concentration of FeIIIL. Equation 18 means that the term a0 ln{a0/(a0 − x)} = kat should be linear with the reaction time t, and the slope of the straight line gives the apparent reaction rate constant ka defined by eq 17. According to eq 16, at the limit of t → 0, the slope in the time course of [FeII] (i.e., the reduction rate) becomes infinity. As x approaches a0, the slope becomes zero. Figures 1−4 (for Zn, Al, and Ti as reducing agents), shown below, seem to
where
K1 =
[Fe IIIL] [Fe3 +][L]
(3)
K
M =2 •M
(4)
where K2 =
[•M] [M]
(5) K
Fe IIIL + •M + H 2O =3 FeHM + OH− + L
(6)
where K3 =
[FeHM][OH−][L] [Fe IIIL][•M] K
FeHM =3 Fe2 + + HM
(7) (8)
where K4 =
[Fe2 +][HM] [FeHM] K
Fe 2 + + L =5 Fe IIL
(9) (10)
where K5 =
[Fe IIL] ([Fe 2 +][L]) k6
HM → M1/ m + + H+
(11) Figure 1. Time courses of the Fe2+ concentration during the reduction of [FeIII(edta)(H2O)]− by aluminum, tin, and zinc at the following conditions: [FeIII(edta)(H2O)]0− = 0.050 M; [EDTA]/[Fe3+] = 1.1; [metal]0/[Fe3+]0 = 10; pH0 = 7; temp. = 20 °C.
(12)
When the rate-determining step is assumed to be reaction 12, the rate of reduction can be written as 14235
dx.doi.org/10.1021/ie502100h | Ind. Eng. Chem. Res. 2014, 53, 14234−14240
Industrial & Engineering Chemistry Research
Article
exhibit the slope of infinity at the limit of t → 0. Equation 18 (i.e., integral reduction rate) has characteristics that (i) as x approaches zero, kat approaches zero, and (ii) as x approaches a0, kat approaches infinity. The use of iron powder as a reducing agent has a merit that the maximum concentrations of NO3− and SO42− in liquid absorbent can be automatically controlled by crystallization of ferric nitrate (Fe(NO3)3) and ferric sulfate (Fe2(SO4)3). Both Fe(NO3)3(s) and Fe2(SO4)3(s) can be removed as the filter cake in the filtration unit. The filtrate will contain such ion species as free Fe 2+ , free Fe 3+ , [Fe II (edta)(H 2 O)] 2− , [FeIII(edta)(H2O)]−, free EDTA, NO3−, SO42−, and so forth. In the case of Fe, where M = Fe and m = 1/2, similarly, eqs 4, 6, 8, and 12 can be considered to reduce to
Fe = •Fe
(19)
Fe IIIL + •Fe + H 2O = Fe2H + OH− + L
(20)
Fe2H = Fe 2 + + HFe
(21)
HFe → Fe 2 + + H+
(22)
The reagents of EDTA and NH4Fe(SO4)2 were supplied by POCH Co. (Poland). Aluminum and tin powders were provided by Riedel-deHean, and zinc powder was supplied by Fluka. All reagents were of analytical grade. Double-distilled water was employed.
4. RESULTS AND DISCUSSION The effects of four parameters, (i) the initial concentration of [FeIII(edta)(H2O]− (0.010−0.100 mol/L), (ii) the initial
Characteristic features of the reduction rate equation, eq 16, are (1) At the limit of t → 0, the reduction rate approaches infinity because of zero [FeIIL]. (2) As the reduction of FeIIIL by metal powder proceeds, the concentration of FeIIL increases. Consequently, the reduction rate decreases and finally approaches zero because of zero [FeIIIL]. (3) The reduction rate is depressed with increasing pH because of apparent reaction rate constant ka in inverse proportion to [OH−], as written by eq 17.
Figure 2. Time courses of the Fe2+ concentration during the reduction of [FeIII(edta)(H2O)]− by zinc at the following conditions: pH0 = 3, 5, 7, and 9; [FeIII(edta)(H2O)]0− = 0.050 M; [Zn0]0/[FeIIIEDTA]0 = 10/1; temp. ≈ 20 °C.
3. EXPERIMENTAL SECTION A glass reactor of 300 cm3 in volume with a reflux condenser equipped with a thermometer and a pH electrode was used. A resistant wire connected through an autotransformer to the current and coiled on the outer wall of the reactor could heat the reaction solution and keep it at a constant temperature higher than ambient one. A 250 cm3 portion of aqueous [FeIII(edta)(H2O]− solution of a desired concentration was prepared inside of the reactor by mixing preliminarily prepared 0.2 M EDTA and 0.5 M NH4Fe(SO4)2 solutions. A 10% excess of EDTA to ferric ions was employed. Both solutions were always prepared and kept under oxygen-free condition. The solution loaded in the reactor was both intensively agitated and isolated from the atmospheric oxygen by the continuous flow of nitrogen of 99.99% purity (MG Chorzow, Poland). The volumetric flow rate of nitrogen was always maintained at the same rate level of 0.20 m3/h. The reduction process was initiated by adding a specified amount of powdered metal to the reactor filled with an aqueous solution of [FeIII(edta)(H2O]−. Next, the pH value of this solution was quickly adjusted by adding solutions of concentrated H2SO4 and/or concentrated NaOH. A prescribed volume of solution was sampled out at the following periods from the beginning of the reaction during each run; 0.5, 3, 6, 9, 12, 15, 20, 30, 40, 50, and 60 min. The concentration of Fe2+ was determined by the o-phenanthroline colorimetry. A 200 UV−Vis Perkin-Elmer spectrophotometer was used.
Figure 3. Time courses of the Fe2+ concentration during the reduction of [FeIII(edta)(H2O)]− by aluminum at the following conditions: [FeIII(edta)(H2O)]0− = 0.005, 0.010, 0.050, and 0.100 M; [Al0]0/ [FeIII(edta)(H2O)]0− = 10/1; pH0 ≈ 7; temp. ≈ 20 °C.
concentration ratio of metal to FeIII (10−40 for Sn) and (10−100 for Al and Zn), (iii) the initial pH value of the solution (3−9), and (iv) the reaction temperature (20−80 °C), were investigated on the time dependency of the Fe II concentration. Typical examples of the experimental results 14236
dx.doi.org/10.1021/ie502100h | Ind. Eng. Chem. Res. 2014, 53, 14234−14240
Industrial & Engineering Chemistry Research
Article
tration of FeIIL increases. Consequently, the reduction rate decreases and finally approaches zero because of zero [FeIIIL]. In the case of reduction of [FeIII(edta)(H2O]− by Na2S2O4,16 a similar kinetic situation was met. 4.2. Al and Sn Powders. In order to evaluate the magnitude of the reduction rate with Al and Sn powders, the pseudo-first-order reaction rate constant (kobs) was tentatively calculated as a measure of the magnitude of the reduction rate. Typical examples of time dependency of [FeII(edta)(H2O]2− during the reduction of [FeIII(edta)(H2O]− by Al and Sn powders are shown in Figures 3 and 4, respectively. The tentative values of the rate constant (kobs) with Al and Sn are listed in Tables 1 and 2, respectively. In Figure 3, the FeII concentration obtained by Al powder does not increase further as the reaction proceeds for some time. Maybe, the reactor here is operated batch-wise with respect to the liquid phase with suspended metal powders, and accordingly, FeIII is believed to be consumed by the reduction, which results in the suppression of FeIII consumption. From these two tables, the comparison of kobs values with Al and Sn indicates that the kobs values with Al are about 10 times larger than those with Sn. The influence of the initial pH value of the solution on the reduction rate is rather complex for Al and Sn. In the case of Al, only at pH0 7 could the value of kobs tentatively be calculated. For acidic and alkaline conditions, the pseudo-first-order kinetic law does not fit the data, as depicted in Figure 3. In the case of Sn, Table 2 and Figure 4 show that kobs decreases with an increase in pH0, and at pH0 9, the pseudo-first-order kinetic law does not fit the data. Such a trend of the reduction rate conforms to eqs 16 and 17, where the rate is inversely proportional to [OH−]. (See the characteristic feature of the reduction rate equation, eq 16, (3) the reduction rate is depressed with increasing pH because of the apparent reaction rate constant ka in inverse proportion to [OH−], as written by eq 17.) It is shown from the effect of the reaction temperature that Al is a very efficient reducing metal at temperatures higher than 40 °C. 4.3. Working Solution in a Simultaneous NO and SO2 Removal Process as a Solvent of [FeIII(edta)(H2O]− Solution. One of the authors (T.T.S.) put in operation the industrial-scale simultaneous removal of NO and SO2 by a
Figure 4. Time courses of the Fe2+ concentration during the reduction of [FeIII(edta)(H2O)]− by tin at the following conditions: pH0 = 3, 5, 7, and 9; [FeIII(edta)(H2O)]0− = 0.050 M; [Sn0]0/[FeIIIEDTA]0 = 10/ 1; temp. ≈ 20 °C.
with Zn, Al, and Sn as reducing agents are shown in a lump in Figure 1. 4.1. Zn Powder. Figure 1 indicates that the reduction rate with Zn powders is very fast as compared to the reduction with Al and Sn. Figure 2 shows the typical time dependency of the concentration of [FeII(edta)(H2O]2− during the reduction of [FeIII(edta)(H2O]− by zinc powder. The reduction rate with Zn can be regarded as almost instantaneous compared to rates with aluminum and tin powders. The concentration of FeII in each first sample, taken after 30 s from the beginning of the run, was always nearly equal to the initial concentration of [FeIII(edta)(H2O]. It implies that [FeIII(edta)(H2O]− is completely reduced to [FeII(edta)(H2O]2−. In summary, the reduction rate approaches infinity at the limit of t → 0 and approaches zero within 30 s Such a trend of the reduction can be suggested from the reduction rate equation, eq 16. Characteristic features of the reduction rate equation, eq 16, are as follows: (1) At the limit of t → 0, the reduction rate approaches infinity because of zero [FeIIL], and (2) as the reduction of FeIIIL by metal powder proceeds, the concen-
Table 1. Experimental Conditions on the Reduction of [FeIII(edta)(H2O)]− by Al and Tentatively Calculated Values of the Pseudo-First-Order Rate Constant parameter III
[Fe (edta)(H2O)]0
−
[FeIII(edta)(H2O)]0− M
[Al0]0/[FeIII]0
pH0
temp. °C
0.010 0.050 0.100
10
7
20
0.050
10 40 100
7
20
0.050
10
[Al0]0/[FeIII]0
pH0
3 5 7 9
temp. 0.050
a
10
7
20
20 40 60 80
kobs min−1 0.0377 0.0368 0.0308 0.0368 0.1195 0.2432 a a 0.0368 a 0.0368 b b b
± ± ± ± ± ±
0.0062 0.0026 0.0025 0.0026 0.0057 0.0117
± 0.0026 ± 0.0026
Experimental data do not fit first-order kinetic law. bToo high to calculate kobs. 14237
dx.doi.org/10.1021/ie502100h | Ind. Eng. Chem. Res. 2014, 53, 14234−14240
Industrial & Engineering Chemistry Research
Article
Table 2. Experimental Conditions on the Reduction of [FeIII(edta)(H2O)]− by Sn and Tentatively Calculated Values of the Pseudo-First-Order Rate Constant parameter III
[Fe (edta)(H2O)]0
−
[FeIII(edta)(H2O)]0− M
[Sn0]0/[FeIII]0
pH0
temp. °C
0.010 0.050 0.100
10
7
20
0.050
10 20 40
7
20
0.050
10
[Sn0]0/[FeIII]0
pH0
3 5 7 9
temp. 0.050
a
10
7
20
20 40 60 80
kobs min−1 0.0039 0.0046 0.0053 b 0.0067 0.0128 0.0531 0.0328 0.0046 b 0.0046 0.0073 0.0125 0.0131
± 0.0005 ± 0.0006a ± 0.0007 ± ± ± ± ±
0.0009 0.0013 0.0042 0.0026 0.0006a
± ± ± ±
0.0006a 0.0008 0.0024 0.0023
Average value calculated from the other two kobs values of the [FeIII(edta)(H2O)]0− series. bExperimental data do not fit first-order kinetic law.
Figure 6. Time courses of the Fe2+ concentration during the reduction of [FeIII(edta)(H2O)]− by aluminum in real technological solution and redistilled water at the following conditions: [FeIII(edta)(H2O)]0− = 0.050 M; [Al]0/[FeIII]0 = 10/1; pH0 = 8; temp. = 20 °C.
Figure 5. Time courses of the Fe2+ concentration during the reduction of [FeIII(edta)(H2O)]− by zinc in real technological solution and redistilled water at the following conditions: [FeIII(edta)(H2O)]0− = 0.050 M; [Zn]0/[FeIII]0 = 10/1; pH0 = 8; temp. = 20 °C.
solvent, it is apparent from these three figures that the working solution compares to or is rather superior to the redistilled water. Accordingly, the powders of metals such as Zn, Al, and Sn work well as reducing agents in the working solution as well as the redistilled water solution of [FeIII(edta)(H2O]−.
chelate method. The simultaneous removal unit was installed at a District Heating Plant located in Trzebinia town, Poland. The heating plant is hard coal-fired. The working solution was a solvent of the [FeIII(edta)(H2O]− solution instead of doubledistilled water and contained the following anion and cation ions in mg/dm3: Cl−, 27000; CO32−, 600; F−, 21.8; NO2−, 35; NO3−, 18; SO42−, 5500; Ca2+, 11500; Co2+, 0.45; Cr3+, 0.4; Cu2+, 0.1; Fe2+, 5.3; Fe3+, 2.5; K+, 600; Mg2+, 1000; Mn2+, 12; Na+, 1800; Ni2+, 0.85; Pb2+, 1.22; and EDTA4−, 8.58. The initial pH value of this solution was 2.57. Afterward, this solution was diluted two-fold with redistilled water for the present laboratory test. For the sake of comparison, redistilled water was used as a solvent of the [FeIII(edta)(H2O]− solution instead of the working solution. Figures 5−7 indicate typical examples of the experimental results on the reduction of [FeIII(edta)(H2O]− by Zn, Al, and Sn powders, respectively, using a working solution in the simultaneous removal process and redistilled water as a solvent of the [FeIII(edta)(H2O]− solution. As to the reactivity of the
5. CONCLUSION The reduction kinetics of ferric ions by metal in aqueous solutions of [FeIII(edta)(H2O]− was investigated under the conditions (1) [FeIII(edta)(H2O]0− = 0.100−0.010 M, (2) [metal]0/[FeIII(edta)(H2O)]0− = 10−100, (3) pH0 = 3−9, and (4) temperature = 20−80 °C. The following findings could be obtained: (1) Zn can effectively act as a reducing agent to [FeIII(edta)(H2O]− even at 20 °C. The reduction rate can be regarded as almost instantaneous compared to rates with aluminum and tin powders. (2) The reduction of the ferric ion by Al proceeds with a very high rate at temperatures higher than 40 °C. 14238
dx.doi.org/10.1021/ie502100h | Ind. Eng. Chem. Res. 2014, 53, 14234−14240
Industrial & Engineering Chemistry Research
Article
(10) Maigut, J.; Meier, R.; van Eldik, R. Influence of Fluoride on the Reversible Binding of NO by [FeII(EDTA)(H2O)]2−. Inhibition of Autoxidation of [FeII(EDTA)(H2O)]2−. Inorg. Chem. 2008, 47, 6314. (11) Wubs, H. J.; Beenackers, A. A. C. M. Kinetics of the oxidation of ferrous chelates of EDTA and HEDTA in aqueous solution. Ind. Eng. Chem. Res. 1993, 32, 2580. (12) Gambardella, F.; Ganzeveld, I. J.; Winkelman, J. G. M.; Heeres, E. J. Kinetics of the Reaction of FeII(EDTA) with Oxygen in Aqueous Solutions. Ind. Eng. Chem. Res. 2005, 44, 8190. (13) Belanzoni, P.; Bernasoni, L.; Baerends, E. J. O2 Activation in a Dinuclear FeII(EDTA) Complex: Spin Surface Crossing as a Route to Highly Reactive Fe(IV)oxo Species. J. Phys.Chem. A 2009, 43, 11926. (14) Bernasoni, L.; Baerends, E. J. Generation of Ferryl Species through Dioxygen Activation in Iron/EDTA Systems: A Computational Study. Inorg. Chem. 2009, 48, 527. (15) Seibig, S.; van Eldik, R. Kinetics of [FeII(edta)] Oxidation by Molecular Oxygen Revisited. New Evidence for a Multistep Mechanism. Inorg. Chem. 1997, 36, 4115. (16) Suchecki, T. T.; Mathews, B.; Kumazawa, H. Kinetic Study of Ambient-Temperature Reduction of FeIIIedta by Na2S2O4. Ind. Eng. Chem. Res. 2005, 44, 4249. (17) Suchecki, T. T.; Kumazawa, H. Application of Hydrazine to Regeneration of Post-Absorption Solutions in Combined SO2/NOx Removal from Flue Gases by a Chelate Method. Sep. Technol. 1994, 763−770. (18) Mendelsohn, M. H.; Harkness, J. B. L. Enhanced flue-gas denitrification using ferrous-EDTA and a polyphenolic compound in an aqueous scrubber system. Energy Fuel. 1991, 5, 244. (19) van der Maas, P.; van den Brink, P.; Klapwijk, B.; Lens, P. Acceleration of the Fe(III)EDTA− reduction rate in BioDeNOx reactors by dosing electron mediating compounds. Chemosphere 2009, 75, 243. (20) van der Maas, P.; Manconi, I.; Klapwijk, B.; Lens, P. Nitric oxide reduction in BioDeNOx reactors: kinetics and mechanism. Biotechnol. Bioeng. 2008, 100, 1099. (21) Kumaraswamy, R.; Kuenen, J. G.; Kleerebezem, R.; van Loosdrecht, M. C.; Muyzer, G. Structure of microbial communities performing the simultaneous reduction of Fe(II)EDTA·NO2− and Fe(III)EDTA−. Appl. Microbiol. Biotechnol. 2006, 73, 922. (22) van der Maas, P.; van den Brink, P.; Utomo, S.; Klapwijk, B.; Lens, P. NO removal in continuous BioDeNOx reactors: Fe(II). Biotechnol. Bioeng. 2006, 94, 575. (23) Kumaraswamy, R.; van, D. U.; Kuenen, J. G.; Abma, W.; van Loosdrecht, M. C.; Muyzer, G. Characterization of microbial communities removing nitrogen oxides from flue gas: the BioDeNOx process. Appl. Environ. Microbiol. 2005, 71, 6345. (24) Kumaraswamy, R.; Muyzer, G.; Kuenen, J. G.; Loosdrecht, M. C. Biological removal of NOx from flue gas. Water Sci. Technol. 2004, 50, 9. (25) van der Maas, P.; van de Sandt, T.; Klapwijk, B.; Lens, P. Biological reduction of nitric oxide in aqueous Fe(II)EDTA solutions. Biotechnol. Prog. 2003, 19, 1323. (26) Zhu, H.; Mao, Y.; Yang, X.; Chen, Y.; Long, X.; Yuan, W. Simultaneous absorption of NO and SO2 into Fe(II)−EDTA solution coupled with the Fe(II)−EDTA regeneration catalyzed by activated carbon. Sep. Purif. Technol. 2010, 74, 1. (27) Yang, X.; Yang, L.; Dong, L.; Long, X.; Yuan, W. Kinetics of the [Fe(III)-EDTA]− Reduction by Sulfite under the Catalysis of Activated Carbon. Energy Fuels 2011, 25, 4248. (28) Long, X.-l.; Yang, L.; Yang; Chou, X.-w.; Li, C.; Yuan, W.-k. Reduction of [Fe(III)EDTA]− Catalyzed by Activated Carbon Modified With Ammonia Solution. Environ. Prog. Sustainable Energy 2014, 33, 99. (29) Wang, L.; Zhao, W.; Wu, Z. Simultaneous absorption of NO and SO2 by FeIIEDTA combined with Na2SO3 solution. Chem. Eng. J. 2007, 132, 227. (30) Manconi, I.; van der Maas, P.; Lens, P. N. L. Effect of sulfur compounds on biological reduction of nitric oxide in aqueous Fe(II)EDTA2− solutions. Nitric Oxide 2006, 15, 40.
Figure 7. Time courses of the Fe2+ concentration during the reduction of [FeIII(edta)(H2O)]− by tin in real technological solution and redistilled water at the following conditions: [FeIII(edta)(H2O)]0− = 0.050 M; [Sn]0/[FeIII]0 = 10/1; pH0 = 8; temp. = 20 °C.
(3) Sn is not recommended as a reducing agent of the ferric ion in the present chelate system.
■
AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. Tel: +48-71-320 38 22. Fax: +48-71-320 35 32. Notes
The authors declare no competing financial interest. ⊥ E-mail:
[email protected] (B.M.); tomasz.suchecki@pwr. edu.pl (T.T.S.);
[email protected] (H.K.).
■
REFERENCES
(1) Krzyzynska, R.; Hutson, N. D. Effect of solution pH on SO2, NO(x), and Hg removal from simulated coal combustion flue gas in an oxidant-enhanced wet scrubber. J. Air Waste Manage. Assoc. 2012, 62, 212. (2) Chien, T. W. Tsung; Hsueh, H. T.; Chu, B. Y.; Chu, H. Absorption kinetics of NO from simulated flue gas using Fe(II)EDTA solutions. Process Saf. Environ. Prot. 2009, 87, 300. (3) Augustyniak, A. W.; Suchecki, T. T. A chelate method for NOx removal from flue gases. State-of-the-art and future. Przem.Chem. 2012, 12, 107. (4) Begel, S.; van Eldik, R. Thermodynamic and kinetic studies on the interaction of RuIII(edta) with NO in an ionic liquid. Dalton Trans. 2011, 40, 4892. (5) Wang, X.; Zhang, Y.; Dong, X.; Chen, M.; Shi, Z.; Zhou, J. Fe(II)EDTA-NO Reduction by Sulfide in the Anaerobic Aqueous Phase: Stoichiometry and Kinetics. Energy Fuels 2013, 27, 6024. (6) Gambardella, F.; Alberts, M. S.; Winkelman, J. G. M.; Heeres, E. J. Experimental and Modeling Studies on the Absorption of NO in Aqueous Ferrous EDTA Solutions. Eng. Chem. Res. 2005, 44, 4234. (7) Du, Q.; Haoyang, Z.; Leiling, M.; Jianmin, G.; Guangbo, Z.; Shaohua, W. Experimental Investigation on the Kinetics of NO Complex Absorption through FeIIEDTA Solvent in a Double-Stirred Reactor. Ind. Eng. Chem. Res. 2011, 50, 4425. (8) Wu, Z. B.; Wang, L.; Zhao, W. R. Kinetic study on regeneration of FeIIEDTA in the wet process of NO removal. J. Chem. Eng. 2008, 140, 130. (9) Liu, Y. X.; Zhang, J.; Wang, Z. L. A study on kinetics of NO absorption from flue gas by using UV/Fenton wet scrubbing. Chem. Engin. J. 2012, 197, 468. 14239
dx.doi.org/10.1021/ie502100h | Ind. Eng. Chem. Res. 2014, 53, 14234−14240
Industrial & Engineering Chemistry Research
Article
(31) Wang, X.; Zhang, Y.; Dong, X.; Chen, M.; Shi, Z.; Zhou, J. Fe(II)EDTA−NO Reduction by Sulfide in the Anaerobic Aqueous Phase: Stoichiometry and Kinetics. Energy Fuels 2013, 27, 6024. (32) Verma, P. S.; Venkateshwar Rao, T.; Jayaraman, A.; Murthy, P. S. N.; Joshi, G. C. Potentiometric estimation of sulfide ion using the Fe(III)-EDTA complex. J. Anal. Chem. 1994, 348, 742. (33) Limtrakul, S.; Rojanamatin, S.; Vatanatham, T.; Ramachandran, P. A. Gas-Lift Reactor for Hydrogen Sulfide Removal. Ind. Eng. Chem. Res. 2005, 44, 6115.
14240
dx.doi.org/10.1021/ie502100h | Ind. Eng. Chem. Res. 2014, 53, 14234−14240
