Article pubs.acs.org/IECR
Understanding the Catalyst Regeneration Kinetics in the Chelated Iron Dehydrosulfurization Process: A Model in Terms of Fe(II) Speciation Lin-Feng Zhai,* Li-Li Hu, and Min Sun Department of Chemical Engineering, Hefei University of Technology, Hefei 230009, China ABSTRACT: Chelated iron dehydrosulfurization process is among the most promising techniques for sulfur recovery from hydrogen sulfide waste. The oxidation of ferrous iron (Fe(II)) plays a crucial role in determining the efficiency of chelated iron process. This work aims to develop a model that describes the electrochemical and aerobic oxidation kinetics of Fe(II) in terms of Fe(II) speciation. By using such a kinetic model, the oxidation behavior of Fe(II) as a function of solution pH and solution composition is appropriately interpreted in the chelated iron process. The oxidation rate of Fe(II) in the EDTA-chelated iron process is mainly controlled by the Fe(EDTA)2− and Fe(OH)(EDTA)3−species, wheras the Fe(NTA)− and Fe(NTA)24− species are the dominant species responsible for the oxidation kinetics of Fe(II) in NTA-chelated iron process. It is anticipated that the kinetic model developed in this work will provide valuable information for a better understanding and manipulation of the chelated iron process.
1. INTRODUCTION Hydrogen sulfide (H2S) is a well-known toxic gas that is produced from petroleum refining, pulp manufacturing, natural gas production, and aerobic and anaerobic wastewater treatments. So far, a variety of approaches have been developed to remove H2S from gaseous streams, among which the chelated iron dehydrosulfurization process has been receiving intensive attention because of its high H2S removal efficiency, great operation flexibility, and good economic potential.1 The chelated iron process is a typical liquid oxidation process based on the ferric (Fe(III))/ferrous (Fe(II)) redox couple. In this process, the Fe(III) is held in solution by chelating agents, such as ethylenediamine tetraacetic acid (EDTA) or nitrilotriacetic acid (NTA). The reaction intent is to absorb the H2S into the chelated Fe(III) solution and oxidize it to elemental sulfur. Meanwhile, the chelated Fe(III) is reduced to the chelated Fe(II) state, which is subsequently oxidized in an aerobic reactor or an air-cathode fuel cell to regenerate chelated Fe(III).2,3 Because the Fe(III)/Fe(II) couple is recycled in the whole process, it is regarded as a redox catalyst in the reaction of H2S with oxygen (O2).4 Compared with the fast kinetics of H2S adsorption reaction, the catalyst regeneration reaction is relatively slower and thereby is the limiting step in the chelated iron process. A large amount of effort has been undertaken to illuminate the aerobic oxidation kinetics of chelated Fe(II).5−8 These investigations have confirmed that the aerobic oxidation rate of Fe(II) chelated with EDTA or similar chelating agents is many orders of magnitude faster than that of free Fe(II) ion (Fe2+).7 However, controversy is encountered when interpreting the kinetic data obtained from different process studies. It is generally accepted that the aerobic oxidation of Fe(II) is firstorder with respect to the concentration of O2, whereas the order of reaction with respect to the concentration of chelated Fe(II) varies between one and two with different operation parameters.5,6,8 What’s more, because the air-cathode fuel cell © 2015 American Chemical Society
technology has only recently been incorporated into the chelate iron process,3 information on the electrochemical oxidation kinetics of chelated Fe(II) is very sparse. For a better manipulation of the catalyst regeneration reaction, it is necessary to establish a unified model that is able to explain the oxidation kinetics of chelated Fe(II) with variable operating parameters. Owing to the importance of iron cycling on natural aquatic systems, over the years, a large number of studies have been undertaken to elucidate the Fe(II) oxidation kinetics in freshwater and saltwater environments.9−19 These studies have shown convincingly that the chemical oxidation of Fe(II) to Fe(III) is a complex process involving a variety of Fe(II) species with different kinetic activities. The Fe(II) oxidation with O2 in natural waters follows the Haber−Weiss mechanism, with reactions 1 and 3 being rate-limiting:9 Fe(II) + O2 → Fe(III) + O2•−
(1)
2H+
Fe(II) + O2•− ⎯⎯⎯→ Fe(III) + H 2O2
(2)
Fe(II) + H 2O2 → Fe(III) + OH• + OH−
(3)
Fe(II) + OH• → Fe(III) + OH−
(4)
When the Fe(II) concentration is at micromolar levels, the steady-state concentrations of O2·−, H2O2, and OH· are reached rapidly in all the kinetic runs, and a 4:1 stoichiometry of Fe(II) oxidation by O2 is expected. The overall Fe(II) oxidation rate, expressed as an apparent oxidation rate that is independent of the mechanism describing the process, can be interpreted as the weighted sum of the oxidation rates of Received: Revised: Accepted: Published: 6430
February 20, 2015 May 3, 2015 May 14, 2015 May 14, 2015 DOI: 10.1021/acs.iecr.5b00716 Ind. Eng. Chem. Res. 2015, 54, 6430−6437
Article
Industrial & Engineering Chemistry Research
air-cathode fuel cell was a 175 mL single-chamber reactor with a 4 × 2.5 cm2 carbon paper (non-wet-proofed, 090S, Toray Co., Japan) as anode and a 2× 2 cm2 Pt-coated carbon paper as the cathode.21 The anodic chamber was filled with the chelating solution and then purged with a mixture of CO2/N2 for 20 min. The pH was adjusted under continuous CO2/N2 sparging, and the FeSO4·7H2O was subsequently added to the anodic chamber at 5 mM as the source of Fe(II). The fuel cell was started by connecting the anode and cathode electrodes with a 1 kΩ resistor. The aerobic reactor was a 250 mL glass cylinder. The chelating solution was first bubbled by air for 20 min, and the solution pH was adjusted. The 5 mM FeSO4·7H2O was then added and oxidized with continuous air bubbling. The experiments were conducted in duplicate at an ambient temperature of 25 °C. 2.4. Analytical Methods. All pH measurements were performed on a FE20K Mettler-Toledo pH meter combined with an LE438 pH electrode. The pH electrode was calibrated using NIST buffer solutions (pH 4.01, 6.86, and 9.18, Leici Co. China). Total Fe(II) concentrations were determined spectrophotometrically using the modified phenanthroline method.22 Absorbance of the Fe(II)−phenanthroline complex was measured at 511 nm on a UV spectrophotometer (UV2401PC; Shimadzu Co., Japan). 2.5. Modeling the Oxidation Kinetics of Micromolar Fe(II) in the Presence of Chelating Agents. The oxidation kinetics of micromolar Fe(II) at any given pH can be described by a general equation as23
individual Fe(II) species. Clearly, it is important to fully characterize the speciation of Fe(II) if one desires to adequately understand the Fe(II) oxidation kinetics. Because the micromolar Fe(II) follows a pseudo-first-order kinetics in the oxidation, the kinetic study is usually performed at such a concentration level as to facilitate the quantification of oxidation activities for individual Fe(II) species.15 This work aims to develop a unified kinetics model that adequately describes the electrochemical and aerobic oxidation behavior of chelated Fe(II). The Fe(II) oxidation kinetics are examined at micromolar concentrations of Fe(II), which is chelated by EDTA or NTA. The kinetics activities of individual chelated Fe(II) species are quantified, and the main species responsible for the overall oxidation rate of Fe(II) are clarified. By using such a model, the oxidation behavior of Fe(II) as a function of solution pH and solution composition is appropriately predicted in the chelated iron process.
2. EXPERIMENTAL SECTION 2.1. Chemicals. All chemicals were purchased from SigmaAldrich and of analytical grade. The nitrogen (N2) and carbon dioxide (CO2) gases were purchased from ZongYi Chemical Co., China and had a purity of 99.999%. The NTA and sodium salt of EDTA were supplied as chelating agents. The 50 mM Fe(II) stock solution was prepared by dissolving FeSO4·7H2O in 2 mM HCl. This level of acidity was sufficient to prevent Fe(II) oxidation over time, yet did not induce significant pH change when the Fe(II) stocks were added into the sample solutions. All solutions were prepared using 18 MΩ Milli-Q water. 2.2. Experiments on the Oxidation of Micromolar Fe(II). The electrochemical oxidation of micromolar Fe(II) was studied in a 250 mL three-electrode electrolysis cell with a glassy carbon working electrode (3 mm in diameter), Pt counter electrode, and saturated calomel electrode (SCE) as the reference. The electrolyte was 200 mM NaCl that contained 9 mM chelating agents. Prior to experiment, the electrolyte solution was purged with a mixture of CO2 and N2 for 20 min to remove O2, and the solution pH was adjusted by HCl and NaOH under continuous CO2/N2 sparging. Next, the Fe(II) stocks were added into the electrolyte to a final concentration of 30 μM. The experiment was started by imposing a fixed potential of 0.0 V (vs SCE) through a CHI 660D electrochemical workstation (CH Instruments Inc., USA). Such a potential value was chosen to simulate the typical anodic potential of an air-cathode fuel cell.20 The aerobic oxidation of micromolar Fe(II) was carried out in a 250 mL glass reactor. The solution containing 200 mM NaCl and 9 mM chelating agents was saturated with O2 by air bubbling, and then the solution pH was adjusted to the desired value. The addition of 30 μM Fe(II) to the reaction cell corresponded to the zero time of the reaction. All experiments were conducted at an ambient temperature of 25 °C, and the solutions were magnetically stirred at a speed of 1250 rpm throughout the experiment. 2.3. Experiments on the Catalyst Regeneration in Chelated Iron Process. The catalyst regeneration reaction in the chelated iron process was simulated by performing the Fe(II) oxidation in an air-cathode fuel cell and an aerobic reactor, respectively. The experiments were conducted in a chelating solution containing 200 mM NaCl, 50 mM NaHCO3, and EDTA or NTA as chelating agents. The concentration of the chelating agent was 7.5 mM unless otherwise indicated. The
d[Fe(II)]T = −kapp[Fe(II)]T [oxidant] dt
(5)
where [Fe(II)]T denotes total Fe(II) concentration and kapp is an apparent oxidation rate constant. At a fixed applied potential, the electrochemical oxidation of Fe(II) on any certain electrode is assumed to follow a pseudo-first-order kinetics with respect to total Fe(II) as18 d[Fe(II)]T = −k′[Fe(II)]T dt
(6)
where k′ is a pseudo-first-order rate constant that is dependent upon the electrode property and solution composition. Similarly, in the O2-saturated solution where ambient O2 concentration, [O2], is in excess of [Fe(II)]T, the aerobic oxidation kinetics of Fe(II) is also expressed by eq 2 where k′ = kapp [O2].16 Equation 2 can be integrated to [Fe(II)]T, t = [Fe(II)]T,0 exp( −k′t )
(7)
Hence, a plot of ln{[Fe(II)]T, t/[Fe(II)]T.0} versus −t should be a straight line with slope k′. When the oxidation of Fe(II) is described in terms of a number of parallel reactions involving individual Fe(II) species, the k′ is the weighted sum of rate constants of individual reactions as14 k′ =
∑ kFe(II) αFe(II) i
i
i
(8)
where Fe(II)i represents individual Fe(II) species, kFe(II)i represents pseudo-first-order oxidation rate constant of species Fe(II),i and αFe(II)i is the molar fraction of species Fe(II)i in total Fe(II). The speciation of Fe(II) was calculated using MINEQL +4.6,24 with a list of equilibrium reactions and their 6431
DOI: 10.1021/acs.iecr.5b00716 Ind. Eng. Chem. Res. 2015, 54, 6430−6437
Article
Industrial & Engineering Chemistry Research
plus glob optimization algorithm in firstOpt 1.5 (7D-Soft High Technology Inc., China). The iteration was repeated 1000 times, and the convergence criterion was 1.00 × 10−10.
corresponding stability constants given in Table 1. The formation of solid compounds was also considered in the speciation calculation; however, their formation was negligible because of the excessive supply of chelating agents. The αFe(II)i for species Fe(II)i was obtained from the speciation model. The overall rate constant k′ was calculated by substituting the experimental Fe(II) concentration/time pair of data into eq 7. The kFe(II)i were estimated by substituting the k′ and αFe(II)i into eq 8 at a series of pH values. The multiple linear regression was performed using the conjugate-gradient
3. RESULTS AND DISCUSSION 3.1. Elucidating Fe(II) Oxidation Kinetics in the Presence of Chelating Agents. To fully understand the influence of EDTA and NTA on the oxidation kinetics of Fe(II), a micromolar Fe(II) oxidation experiment was conducted to clarify the oxidation activities of individual chelated Fe(II) species. As shown in Figure 1, both the electrochemical and aerobic oxidation of Fe(II) follow a pseudo-first-order behavior under the experimental conditions. The rate constant, k′, varies significantly from 0.0045 min−1 in the pH 6.23 electrochemical system with NTA to 1.32 min−1 in the pH 6.85 aerobic system with EDTA. The electrochemical oxidation of Fe(II) was more than 1 order of magnitude slower than the aerobic oxidation of Fe(II), and the Fe(II) oxidation rate generally increased with the increase of pH. The k′ obtained at different pHs is fitted to a kinetic model described by eq 4 to calculate the rate constants of individual chelated Fe(II) species. The oxidation of Fe(II) is composed of several parallel reactions in which the individual Fe(II) species is oxidized at different rates. As shown in Figure 2, the Fe(II) is almost totally chelated in the presence of excessive chelating agents, and inorganic Fe(II) species such as the FeCl+, FeSO40, Fe(OH)20, and Fe(OH)+ are present at very low concentrations. The Fe(EDTA)2− is the most abundant species in the Fe(II)-EDTA solution, whereas the Fe(II)-NTA solution is dominated by the Fe(NTA)− species. To facilitate calculation, the contributions of Fe2+, FeCl+, and FeSO40 species to the overall Fe(II) oxidation kinetics are ignored because of their inert activities in both the electrochemical and aerobic systems.10,13,15,18,19 The Fe(OH)20 and Fe(OH)+ species are also neglected from eq 4 owing to their low concentrations. As a result, eq 4 is reduced to eqs 9 and 10 for the Fe(II)-EDTA and Fe(II)-NTA systems, respectively.
Table 1. Stability Constants for the Formation of Fe(II) Complexes Considered in the Speciation Model log K, 25 °C
no.
species
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16
H+ + OH− = H2O H+ + CO32− = HCO3− 2H+ + CO32− = H2CO3 H+ + SO42− = HSO4− Fe2+ + H2O = FeOH+ + H+ Fe2++ 2H2O = Fe(OH)20 + 2H+ Fe2+ + CO32− = FeCO30 Fe2+ + H+ + CO32− = FeHCO3+ Fe2+ + 2CO32− = Fe(CO3)22− Fe2+ + H2O + CO32− = Fe(OH)CO3− + H+ Fe2+ + Cl− = FeCl+ Fe2+ + SO42− = FeSO40 Na+ + CO32− = NaCO3− Na+ + H+ + CO32− = NaHCO3 Na+ + SO42− = NaSO4− H+ + EDTA4− = H(EDTA)3−
17
2H+ + EDTA4− = H2(EDTA)2−
17.22
18
3H+ + EDTA4− = H3(EDTA)−
20.34
19
4H+ + EDTA4− = H4(EDTA)0
22.50
20
5H+ + EDTA4− = H5(EDTA)
+
24.00
21
Fe2+ + EDTA4− = Fe(EDTA)
2−
16.00
22
24
Fe2+ +2OH− + EDTA4− = Fe(OH)2(EDTA)4− Fe2+ +OH− + EDTA4− = Fe(OH) (EDTA)3− Fe2+ + H+ + EDTA4− = FeH(EDTA) −
25
Na+ + EDTA4− = Na(EDTA)3−
26
H+ + NTA3− = H(NTA)2−
27
2H+ + NTA3− = H2(NTA)−
23
+
3−
14.00 10.30 16.70 1.99 −9.51 −20.60 5.69 11.80 7.45 −4.03 0.30 2.42 1.27 10.10 1.06 10.95
−4.00 6.50 19.06 2.70 10.28
0
13.22
28
3H + NTA
29
4 H+ + NTA3− = H4(NTA)+
16.22
30
Fe2+ + NTA3− = Fe(NTA)−
10.19
31
Fe2+ + 2NTA3− = Fe(NTA)24−
12.62
32
Fe2+ + OH− + NTA3− = Fe(OH) (NTA)2− Fe2+ + H+ + NTA3− = FeH(NTA)0
−1.06
33
= H3(NTA)
15.22
12.29
reference 25 25 25 26 26 26 10 27 10 10 10 10 10 10 10 supplied by Mineql supplied by Mineql supplied by Mineql supplied by Mineql supplied by Mineql supplied by Mineql supplied by Mineql supplied by Mineql supplied by Mineql supplied by Mineql supplied by Mineql supplied by Mineql supplied by Mineql supplied by Mineql supplied by Mineql supplied by Mineql supplied by Mineql supplied by Mineql
k′ = kFe(EDTA)2−αFe(EDTA)2− + kFe(OH)2 (EDTA)4− αFe(OH)2 (EDTA)4− + kFe(OH)(EDTA)3−αFe(OH)(EDTA)3− + kFeH(EDTA)−αFeH(EDTA)−
(9)
k′ = kFe(NTA)−αFe(NTA)− + kFe(NTA)2 4−αFe(NTA)2 4− + kFe(OH)(NTA)2−αFe(OH)(NTA)2− + kFeH(NTA)0αFeH(NTA)0 (10)
By substituting the k′ and αFe(II)i into eqs 9 and 10, rate constants of the eight individual chelated Fe(II) species are estimated in Table 2. As anticipated, the rate constants vary greatly among the different Fe(II) species and generally present higher values for aerobic oxidation than for electrochemical oxidation. The Fe(OH)2(EDTA)4−, FeH(EDTA)−, Fe(OH)(NTA)2−, and FeH(NTA)0 species are identified as inactive species in the oxidation, evidenced by their extremely low rate constants. The Fe(EDTA)2− and Fe(NTA)− species show moderate kinetic reactivity, and the Fe(OH)(EDTA)3− and Fe(NTA)24− are the most kinetically active species in both the aerobic and electrochemical systems. To verify the validity of estimated kinetic characteristics of chelated Fe(II) species, the overall Fe(II) oxidation rate, k′, is predicted by substituting the rate constants of individual 6432
DOI: 10.1021/acs.iecr.5b00716 Ind. Eng. Chem. Res. 2015, 54, 6430−6437
Article
Industrial & Engineering Chemistry Research
Figure 1. Pseudo-first-order kinetics for the (A) electrochemical and (B) aerobic oxidation of 30 μM Fe(II) in 9 mM EDTA solution and (C) electrochemical and (D) aerobic oxidation of 30 μM Fe(II) in 9 mM NTA solution.
Table 2. Estimated Pseudo-First-Order Rate Constants for Individual Chelated Fe (II) Species in the Electrochemical and Aerobic Oxidation pseudo-first-order rate constants species
electrochemical oxidation (min−1)
Fe(EDTA)2− Fe(OH)2(EDTA)4− Fe(OH)(EDTA)3− FeH(EDTA)− Fe(NTA)− Fe(NTA)24− Fe(OH)(NTA)2− FeH(NTA)0
5.97 × 10−3
