Ind. Eng. Chem. Res. 2005, 44, 4249-4253
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Kinetic Study of Ambient-Temperature Reduction of FeIIIedta by Na2S2O4 Tomasz T. Suchecki,*,†,‡ Barbara Mathews,§ and Hidehiro Kumazawa| Faculty of Environmental Engineering, Wroclaw University of Technology, Wybrzeze Wyspianskiego 27, 50-379 Wroclaw, Poland, Institute of Multidisciplinary Research for Advanced Materials, Tohoku University, 2-1-1 Katahira, Aoba-ku, Sendai 980-8577, Japan, Institute of Environmental Engineering, Polish Academy of Sciences, ul. M. Curie-Sklodowskiej 34, 41-819 Zabrze, Poland, and Department of Chemical Process Engineering, Toyama University, Toyama 930-0887, Japan
The chelate method can provide simultaneous control of combustion flue gases SO2/NOx by scrubbing into aqueous solutions of ferrous chelates, e.g., FeIIedta, an ethylenediaminetetraacetic acid. Because of oxidation of FeIIedta to completely inactive in the system FeIIIedta, the absorption solution requires regeneration. A new proposal here for this process is based on the direct reduction of FeIIIedta by a dithionite ion, S2O42-. The rate equation for the reduction of FeIIIedta by dithionite was derived in terms of the rate-determining-step approximation method. A kinetic study of this process was conducted by using a bubbling-type batch reactor. The following initial parameters were examined: FeIIIedta and Na2S2O4 concentrations, pH, and oxygen concentration. It was concluded that dithionite is a promising reducing agent for the regeneration of the absorption solution of the chelate method. The ambient-temperature regeneration process should be operated under the following conditions: pH0 ) 3.5-9.0 and [Na2S2O4]0/[FeIIIedta]0 ) 0.6. Introduction There are a number industrial technologies for separate control of sulfur and nitrogen oxides, SO2/NOx, from combustion flue gases. Nitrogen oxides consist of 95% NO and 5% NO2, approximately, and the first one is practically insoluble in water, contrary to SO2 and NO2. Most of technologies for SO2 abatement are based on wet lime or limestone absorption methods, which produce gypsum as the end product. A selective catalytic reduction of NOx by gaseous ammonia is the most widespread process for deep control of nitrogen oxides in flue gases emitted from big power furnaces. However, the complex protection of the atmosphere against emission of sulfur and nitrogen oxides is achieved when both of the above-mentioned methods are combined in series. Such systems have been working in many Japanese electric power stations fueled by hard coal for the last 20 years. However, there are prosperous, but still not commercialized, technologies of air protection against sulfur and nitrogen oxide emissions from combustion flue gases. One of them is a chelate method, of which chemical, kinetic, and process fundamentals have been extensively investigated for more than 20 years by many research groups from USA, Japan, and Europe. This technology was successfully demonstrated by us on the industrial scale (50 000 m3 flue gas flow) at the District Heating Plant in the town of Trzebinia in southern Poland. Recently published papers1-3 suggest that there is continuous interest in this method. * To whom correspondence should be addressed. Tel.: +48-71-320 38 22. Fax: +48-71-320 35 32. E-mail:
[email protected]. † Wroclaw University of Technology. ‡ Tohoku University. § Polish Academy of Sciences. | Toyama University.
The chelate method can provide simultaneous removal of NOx and SO2 in one absorption step by scrubbing flue gases by aqueous solutions of ferrous chelate, FeIIL, iron(II) ion coordinated to a ligand, L. Only ethylenediaminetetraacetic acid (edta) has been widely investigated as the ligand of ferrous ions in this system.1,2,4-13 Another one, nitrilotriacetic acid (nta), has been employed to a much lesser extent.8,14-16 The following compounds were also tested as ferrous ion ligands: (i) acetyl acetonate,4 (ii) citrate,4 (iii) a polymeric chelate,17 (iv) dithiocarbamates,18 (v) cysteine,19 (vi) amino acids and peptides,20 (vii) immobilized resins,21 (viii) bis(2,3-dimercapto-1-propanesulfonate),22 and (ix) tiochelate.23 The ferrous chelate enables nitric oxide to form FeII(L)(NO), nitrosyl ferrous chelate, in an aqueous solution. The nitrosyl ferrous chelate reacts in an absorption solution with dissolved SO2 and undergoes further transformations toward gaseous and liquid products. A detailed chemistry of this process is described elsewhere.9 However, during the absorption process, the Fe2+ ions coordinated to a ligand are oxidized to Fe3+ ions,23-26 which have no ability to bind nitrous oxide. Therefore, the regeneration of a spent absorption solution, a backconversion from a FeIIIL chelate to a FeIIL one, became a fundamental part of the method. Until now, only the sulfite ion has been extensively examined as a reducing agent for the regeneration process when edta was employed as the ligand of ferrous ions.9,12,27-34 Hydrazine has also been shown to be an efficient compound for reduction of ferric ions within FeIIIedta complexes.35 However, the results of the abovementioned researches do not allow one to start the commercialization of this method presumably because of inevitable complicated and expensive treatment of spent absorption solutions. It should be noted that hydrogen sulfite is a very efficient reducing agent of ferric edta chelate in the systems of desulfurization of natural gases.36 However, there were no proposals to
10.1021/ie0493006 CCC: $30.25 © 2005 American Chemical Society Published on Web 05/10/2005
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Ind. Eng. Chem. Res., Vol. 44, No. 12, 2005
use it for the regeneration process, probably because of its high toxic properties. Dithionite ion, S2O42-, is a well-known strong reductant because it consists of two sulfoxyl radicals, •SO2-. It was added as a redox mediator to the FeIIedta solution employed for NO scrubbing and was regenerated by electrochemical reduction of SO32- ions.37 However, this process seems to be very complicated. In this work, the mechanism and kinetics of FeIIIL reduction by bisulfite, HSO3-, are discussed because of the fact that sulfites are always present in the SO2/NOx absorption solution. Next, the stoichiometry of the dithionite ion reaction with FeIIIL is presented. A kinetic equation of this reduction process is developed using an analogy to the sulfite one. Results of detailed kinetic investigations of the reaction between FeIIIedta and dithionite are presented. Finally, process parameters of the absorption solution regeneration are proposed. Kinetic Background
K9 ) [FeIIL]/([Fe2+][L]) k11
•
SO3- + H+ 98 HSO3 k12
HSO3 + HSO3 98 S2O62- + 2H+
(10) (11) (12)
When the rate-determining step is assumed to be reaction (11), the rate of reduction can be given as
-d[FeIIIL]/dt ) k11[•SO3-][H+]
(13)
Elimination of the concentrations of •SO3- and H+ with the help of eqs 6, 8, and 10 yields
-d[FeIIIL]/dt ) k11K5K7K9[HSO3-][FeIIIL]/[FeIIL] ) k14[HSO3-][FeIIIL]/[FeIIL] (14) where
Reduction of FeIIIL by Sulfite. Because sulfite has been intensively examined as the reducing agent to ferric chelates, first this process will be described below in detail. The reduction rate expression will be derived on the basis of a rate-determining-step approximation method. The reduction of Fe3+ by sulfite with a coexisting ligand can be written stoichiometrically as
2FeIIIL + 2HSO3- f 2FeIIL + S2O62- + 2H+ (1)
k14 ) k11K5K7K9
(15)
The concentration of FeIIIL approximately equates the total concentration of Fe3+ in the solution, when the last one is equimolar to the concentration of L. Then, eq 14 reduces to
-d[Fe3+]/dt ) k14[HSO3-][Fe3+]/[FeIIL]
(16)
The reaction has been found to be first-order in FeIIIL and of order -1 in FeIIL for edta38 and nta,39 viz.
When the amount of sulfite consumed by the reduction and coordination of Fe3+ is very small compared to the initial amount of sulfite, [HSO3-]0, eq 16 can be rewritten as
-d[FeIIIL]/dt ) k1[FeIIIL]/[FeIIL]
-d[Fe3+]/dt ) k17[Fe3+]/[FeIIL]
(17)
k17 ) k14[HSO3-]0
(18)
(2)
This type of rate equation can be derived by assuming the following six elementary steps involving a ratedetermining step:38 K3
Fe3+ + L ) FeIIIL
(3)
K3 ) [FeIIIL]/([Fe3+][L])
(4)
where
K5
FeIIIL + HSO3- ) FeSO3+ + H+ + L
where
The reduction rate for ferric ions by sulfite, however, was found to be rather low at temperatures below 100 °C.31 Reduction of FeIIIedta by Dithionite. Dithionite ions dissociate to form two sulfoxyl radicals in aqueous solutions:
(5)
S2O42- T 2•SO2-
(19)
K19 ) [•SO2-]2/[S2O42-]
(20)
where where +
+
III
-
K5 ) [FeSO3 ][H ][L]/([Fe L][HSO3 ]) K7
FeSO3+ ) Fe2+ + •SO3-
(6) (7)
Then, plausible elementary reaction steps of FeIIIL reduction by dithionite ion can be written as
FeIIIL + •SO2- + H2O f FeIIIL + HSO3- +
where
K21
H+ ) FeSO3+ + 2H+ + L (21) K7 ) [Fe2+][•SO3-]/[FeSO3+] K9
Fe2+ + L ) FeIIL where
(8) (9)
where
K21 ) [FeSO3+][H+]2[L]/([FeIIIL][•SO2-]) The combination of eq 22 with eq 20 yields
(22)
Ind. Eng. Chem. Res., Vol. 44, No. 12, 2005 4251
K21 ) [FeSO3+][H+]2[L]/(K191/2[FeIIIL][S2O42-]1/2) (23) Similarly, the rate-determining step is assumed to be reaction (11). By eliminating the concentration of •SO3with the aid of eqs 8, 10, and 23, one gets
-d[Fe3+]/dt ) k24[FeIIIL][S2O42-]1/2/([FeIIL][H+]) (24) where
k24 )
k11K7K191/2K21
2S2O4
-
series no.
[Na2S2O4]0, [FeIIIedta]0, mmol/L mmol/L
parameter
+ H2O T 2HSO3 + S2O3
2-
50
∼3.0
N2
∼7.0
N2
15
5 10 50 100 50
N2
15
50
∼3.5 ∼7.0 ∼9.0 ∼7.0
1
[Na2S2O4]0
2
[FeIIIedta]0
15 25 30 35 15
3
pH0
4
bubbling gas
(26)
in aqueous solutions. When oxygen is available, dithionite can also react according to following reaction:
S2O42- + O2 + 2OH- ) SO42- + SO32- + H2O (27) and the regeneration kinetics becomes more complex. Experimental Section The experimental setup employed in this study was very simple. It consisted of a 250-cm3-volume cylindrical glass reactor equipped with an Alliahn condenser, a pH electrode, and a gas inlet at the reactor bottom. An aqueous solution of FeIIIedta was prepared inside the reactor by mixing edta‚2Na and NH4Fe(SO4)2 solutions. The reduction process was initiated by adding an appropriate volume of the dithionite solution to the reactor. Next, the pH value was quickly adjusted by solutions of H2SO4 and/or alkalized edta. During each run, 1-5-cm3 volumes of the reaction solution were periodically sampled out to check the ferrous ion concentrations by o-phenanthroline colorimetry.40 This analytical method is based on the fact that ferrous ion forms with colorless o-phenanthroline ions a red chelate complex, which is very stable within the pH ) 2.5-9 range. The sample was added to a 5-cm3 volume of a 1.5 wt % solution of o-phenanthroline. Next, 10 cm3 of a 50% solution of an ammonia acetate buffer (pH ) 5.6) and 1 cm3 of an edta solution were added and diluted to 100 cm3 with distilled water. The absorbance of the sample was measured at λ ) 512 nm after a minimum of 1 h. A Perkin-Elmer Lambda 20 spectrophotometer was used. The solutions loaded in the reactor were very intensively mixed and always isolated from the atmospheric oxygen by the continuous flow of nitrogen of 0.18 m3/h. When influences of air or pure oxygen were investigated, the nitrogen gas flow was switched to another one, together with introduction of a dithionite solution to the reactor. Mass flow controllers were employed to maintain the stable gas flows. All experiments were conducted at room temperature, approximately 20 °C. All solutions used in this study were prepared from oxygen-free distilled water. Fresh solutions of edta‚2Na, NH4Fe(SO4)2, and Na2S2O4 were prepared before each run. The solutions were always kept under oxygen-free
a
bubbling gas
pH0
(25)
It should be noted in eq 24 that the concentration of H+ appears on the right-hand side of the denominator. Dithionite ions undergo a disproportionation reaction 2-
Table 1. Initial Experimental Conditions Employed at Room Temperature
N2 N2a air O2
Oxygen-saturated solution.
Table 2. Experimental Results on Time Courses of the Fe2+ Concentration (mmol/L) and Solution pH Values for the No. 1 Series [Na2S2O4]0, mmol/L 15 time, min [Fe2+] 0.5 3 6 9 12 15
23.9 24.0 23.9 23.9 23.7 23.8
25
30
35
pH
[Fe2+]
pH
[Fe2+]
pH
[Fe2+]
pH
3.5 1 3.52 3.55 3.59 3.62 3.66
41.9 41.3 41.2 41.5 41.7 41.2
2.76 2.82 2.92 3.01 3.07 3.15
49.9 49.8 49.7 49.8 49.9 49.8
2.72 2.78 2.86 2.95 3.03 3.11
49.1 49.9 50.6 49.6 50.3 49.5
2.75 2.82 2.89 2.94 3.02 3.11
Table 3. Experimental Results on Time Courses of the Fe2+ Concentration (mmol/L) and Solution pH Values for the No. 2 Series [FeIIIedta]0, mmol/L 5 time, min
[Fe2+]
0.5 3 6 9 12 15
2.3 2.2 2.2 2.3 2.3 2.2
10 pH
[Fe2+]
6.75 6.74 6.76 6.77 6.77 6.78
5.2 5.2 5.1 5.2 5.1 5.2
50 pH
[Fe2+]
6.53 6.55 6.55 6.56 6.56 6.57
24.3 24.8 25.0 24.7 24.5 24.7
100 pH
[Fe2+]
pH
6.97 6.96 6.96 6.96 6.96 6.96
47.7 47.3 47.9 47.8 47.5 47.2
7.09 7.08 7.08 7.08 7.08 7.07
conditions. Sodium dithionite was provided by Fluka, and other chemicals were provided by POCH Co. (Poland). All reagents employed were of analytical grade. Nitrogen and oxygen, both of 99.99% purity, were supplied by MG Chorzow Co. (Poland). Experimental Results and Discussion The effects of four parameters, (1) the initial concentration of Na2S2O4, (2) the initial concentration of FeIIIedta, (3) the initial pH, and (4) the oxygen concentration in the bubbling gas, were investigated on the time dependency of the Fe2+ concentration. All of the experimental conditions are summarized in Table 1. The experimental results in the form of time dependencies of the Fe2+ concentration and pH of the absorption solution are summarized in Tables 2-5. First, the reduction by dithionite is shown to be very fast as compared to the reduction by sulfite. The percentage of Fe3+ ions converted to Fe2+ ones depends generally on the ratio of Na2S2O4 to FeIIIedta at the beginning of each experimental run. The experimental results at different initial concentrations of dithionite,
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Table 4. Experimental Results on Time Courses of the Fe2+ Concentration (mmol/L) and Solution pH Values for the No. 3 Series pH0 ∼3.5 time, min
[Fe2+]
0.5 3 6 9 12 15
23.9 24.0 23.9 23.9 23.7 23.8
∼7.0 pH
[Fe2+]
3.51 3.52 3.55 3.59 3.62 3.66
24.3 24.8 25.0 24.7 24.5 24.7
∼9.0 pH
[Fe2+]
pH
6.97 6.96 6.96 6.96 6.96 6.96
25.0 24.5 24.3 24.8 24.3 24.4
8.86 8.83 8.79 8.77 8.76 8.75
Table 5. Experimental Results on Time Courses of the Fe2+ Concentration (mmol/L) and Solution pH Values for the No. 4 Series N2 time, min [Fe2+] 0.5 3 6 9 12 15 a
24.3 24.8 25.0 24.7 24.5 24.7
N2
The reduction of FeIIIedta by dithionite in aqueous solutions was studied using a bubbling-type batch reactor. The following findings could be obtained. (1) Dithionite can effectively act as a reducing agent to FeIIIedta at room temperature and a wide range of process parameters. (2) The reduction rate can be regarded as almost instantaneous for all experimental conditions covered. (3) The derived reduction rate equation can simulate very high reduction rates observed here. (4) Recommended regeneration conditions for ambient-temperature spent absorption solutions are as follows: pH0 ) 3.5-9.0 and [Na2S2O4]0/[FeIIIedta]0 ) 0.6. Acknowledgment
bubbling gas a
Conclusion
air
O2
pH
[Fe2+]
pH
[Fe2+]
pH
[Fe2+]
pH
6.97 6.96 6.96 6.96 6.96 6.96
20.6 20.4 21.1 20.6 20.5 20.2
7.15 7.08 7.08 7.07 7.06 7.06
25.8 19.7 14.3 8.1 2.2 0.5
6.94 7.05 7.10 7.18 7.22 7.22
10.4 3.1 0.4 0.1 0.1 0.1
7.16 7.25 7.25 7.25 7.25 7.25
Oxygen-saturated solution.
viz., different initial ratios of dithionite to ferric ion, which are listed in Table 2, showed that the reduction efficiency depended on this ratio within the interval of 0.1-0.5. Above the value of 0.5, the whole FeIIIedta was reduced to FeIIedta. Additionally, increases of the solution pH were observed during all experimental runs. Table 3 shows the experimental results on the time courses of the Fe2+ concentration and solution pH at different initial concentrations of FeIIIedta. At different initial FeIIIedta concentrations, the degree of iron converted was kept constant at a level of ca. 50%, irrespective of the initial concentration of FeIIIedta. Solution pHs during all runs remained almost unchanged. Table 4 shows the effects of the initial solution pH on the time dependencies of the Fe2+ concentration and solution pH at constant initial concentrations of Na2S2O4 and FeIIIedta of 15 and 50 mmol/L, respectively. It is apparent that the conversion to FeIIedta increases with an increase in the initial pH. Thus far, only pure N2 was used as the sparging gas. Next, the concentration of oxygen in gas or in liquid was varied. For the sake of comparison, three kinds of bubbling gases were employed: pure nitrogen, air, and pure oxygen. Table 5 lists the experimental results on the time dependencies of the Fe2+ concentration and solution pH for different kinds of bubbling gases at constant initial concentrations of Na2S2O4 and FeIIIedta of 15 and 50 mmol/L, respectively, and an initial solution pH of 7. The first experiment can be used as a reference; ferric ions are converted to ferrous ones with ca. 50% efficiency, and no pH changes are observed. An oxygen-saturated reaction solution decreased the reduction efficiency only by 10%. Air had a negligible influence on the reduction process only at the beginning of the experiment. However, when pure oxygen was used, about 20% of ferric ions were reduced in 0.5 min. In the last two runs, ferrous ions were oxidized completely to ferric ones.
This research was supported by the [Polish] State Committee for Scientific Research. We gratefully appreciate the suggestion of Prof. R. van Eldik of Erlangen-Nurnberg University to use dithionite as the reductant. Literature Cited (1) Demmink, J. F.; van Gils, I. C. F.; Beenackers, A. A. C. M. Absorption of Nitric Oxide into Aqueous Solutions of Ferrous Chelates Accompanied by Instantaneous Reaction. Ind. Eng. Chem. Res. 1997, 36, 4914. (2) 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. (3) Long, X.; Xiao, W.; Yuan, W. Removal of Sulfur Dioxide and Nitric Oxide Using Cobalt Ethylenediamine Solution. Ind. Eng. Chem. Res. 2005, 44, 686. (4) Termoto, M.; Hiramine, S. I.; Shimada, Y.; Sugimoto, Y.; Terannishi, H. Absorption of Dilute Monoxide in Aqueous Solutions of Fe(II)-EDTA and Mixed Solutions of Fe(II)-EDTA and Na2SO3. J. Chem. Eng. Jpn. 1978, 11, 450. (5) Sada, E.; Kumazawa, H.; Kudo, I.; Kondo, T. Individual and Simultaneous Absorption of Dilute NO and SO2 in Aqueous Slurries of MgSO3 with Fe(II)-EDTA. Ind. Eng. Chem. Process Des. Dev. 1980, 19, 377. (6) Littlejohn, D.; Chang, S. G. Kinetic Study of Ferrous Nitrosyl Complexes. J. Phys. Chem. 1982, 86, 537. (7) Sada, E.; Kumazawa, H.; Sawada, Y.; Kondo, T. Simultaneous Absorption of Dilute Nitric Oxide and Sulfur Dioxide into Aqueous Slurries of Magnesium Hydroxide with Added Iron(II)EDTA Chelate. Ind. Eng. Chem. Process Des. Dev. 1982, 21, 771. (8) Chang, S. G.; Littlejohn, D.; Lynn, L. Effects of metal chelates on wet flue gas scrubbing chemistry. Environ. Sci. Technol. 1983, 17, 649. (9) Sada, E.; Kumazawa, H.; Takada, Y. Chemical Reactions Accompanying Absorption of NO into Aqueous Mixed Solutions of Fe(II)-EDTA and Na2SO4. Ind. Eng. Chem. Fundam. 1984, 23, 60. (10) Weisweiler, W.; Blumhofer, R.; Westermann, T. Absorption of Nitrogen Monoxide in Aqueous Solutions Containing Sulfite and Transition-Metal Chelates such as Fe(II)-EDTA, Fe(II)-NTA, Co(II)-Trien and Co(II)-Treten. Chem. Eng. Proc. 1986, 20, 155. (11) Huasheng, L.; Wenchi, F. Kinetics of absorption of Nitric Oxide in Aqueous FeII-EDTA solution. Ind. Eng. Chem. Res. 1988, 27, 770. (12) Sada, E.; Kumazawa, H.; Yoshikawa, Y. Simultaneous Removal of NO and SO2 by Absorption into Aqueous Mixed Solutions. AIChE J. 1988, 34, 1215. (13) Yih, S. M.; Lii, C. W. Absorption of NO and SO2 in Fe(II)EDTA Solutions. I. Absorption in a Double Stirred Vessel. Chem. Eng. Commun. 1988, 73, 43. (14) Lin, N.; Littlejohn, D.; Chang, S. G. Thermodynamics and Kinetics of the Coordination of Nitric Oxide to Iron(II) NTA in Aqueous Solutions. Ind. Eng. Chem. Process Des. Dev. 1982, 21, 725. (15) Sada, E.; Kumazawa, H.; Machida, H. Absorption of Dilute Nitric Oxide into Aqueous Solutions of Sodium Sulfite with Added
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(30) Littlejohn, D.; Chang, S. G. Reaction of Ferrous Chelate Nitrosyl Complexes with Sulfite and Bisulfite Ions. Ind. Eng. Chem. Res. 1990, 29, 10. (31) Suchecki, T. T.; Sada, E.; Kumazawa, H. Reduction of FeIIIedta by Sulfite at Boiling Temperature. Ind. Eng. Chem. Res. 1991, 30, 2201. (32) Dellert-Ritter, M.; van Eldik, R. Kinetics and Mechanism of the Complex Formation of Ethylenediaminetetraacetateiron(III) with Sulfur(IV)-oxides in Aqueous Solution. J. Chem. Soc., Dalton Trans. 1992, 1037. (33) Dellert-Ritter, M.; van Eldik, R. Kinetics and Mechanism of the Redox Behaviour of the Ethylenediaminetetraacetateiron(III)-sulfite System in Aqueous Solution. J. Chem. Soc., Dalton Trans. 1992, 1045. (34) Brandt, Ch.; Fabian, I.; van Eldik, R. Kinetics and Mechanism of the Iron(III)-catalyzed Autoxidation of Sulfur(IV) Oxides in Aqueous Solution. Evidence for the Redox Cycling of Iron in the Presence of Oxygen and Modeling of the Overall Reaction Mechanism. Inorg. Chem. 1994, 33, 687. (35) 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. In Separation Technology; Vasant, E. F., Ed.; Elsevier: Amsterdam, The Netherlands, 1994; p 763. (36) 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. (37) Kleifges, K.-H.; Jezeliunas, E.; Juttner, K. Electrochemical Study of Direct and Indirect NO Reduction with Complexing Agents and Redox Mediator. Electrochim. Acta 1997, 42, 2947. (38) Sato, T.; Shimizu, T.; Okabe, T. The Formation of Dithionite by the Reaction of FeIIIedta with Sodium Sulfite. Nippon Kagaku Kaishi 1978, 361-366. (39) Sada, E.; Kumazawa, H.; Machida, H. Absorption of Dilute Nitric Oxide into Aqueous Solutions of Sodium Sulfite with Added Iron(II)NTA and Reduction Kinetics of Iron(III)NTA by Sodium Sulfite. Ind. Eng. Chem. Res. 1987, 26, 2016. (40) Harvey, A. E., Jr.; Smart, J. A.; Amis, E. S. Simultaneous Spectrophotometric Determination of Iron(II) and Total Iron with 1,10-Phenenthroline. Anal. Chem. 1955, 27, 26.
Received for review August 4, 2004 Revised manuscript received March 16, 2005 Accepted April 15, 2005 IE0493006
