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Ind. Eng. Chem. Res. 2009, 48, 10186–10189
Effect of Dioxygen Partial Pressure on Ligand Degradation in Chelated Iron Dehydrosulfurization Processes Stephen A. Bedell* and Clare M. Worley The Dow Chemical Company, 2301 North Brazosport BouleVard, Freeport, Texas 77566
Chelated iron processes such as LO-CAT and SulFeroxSM utilize ferric chelate solutions for the oxidation of hydrogen sulfide to form aqueous suspensions of sulfur. During the ferrous chelate reoxidation, dioxygen reduction intermediates are believed to initiate oxidative degradation reactions of the aminocarboxylate ligands used in the process. This study shows that the same degree of Fe(II) oxidation (regeneration) can be achieved with reduced ligand degradation by decreasing the rate at which O2 is supplied to a reactor. This apparent mass transfer limited degradation is explained in terms of dioxygen reduction intermediate formation and reactions. A model for bubble column regeneration was used to predict local partial pressures of dioxygen for a number of reactor configurations. Such analysis predicts several changes in regenerator design which could substantially reduce the amount of chemical usage in such a process. Changes include reduction of the stoichiometric excess of air, use of multiple reactors, and recycling of spent air. Introduction Chelated iron processes such as LOCAT and SulFeroxSM utilize ferric chelate solutions for the oxidation of hydrogen sulfide to form aqueous suspensions of sulfur. After filtration and reoxidation of the reduced iron chelates, the solution is again ready for contact with sour gas.1 During the ferrous chelate reoxidation, dioxygen reduction intermediates are believed to initiate oxidative degradation reactions of the aminocarboxylate ligands used in the process.2 Experimental Section Dioxygen Dependent Regeneration. A 600 mL jacketed glass reactor (30 cm × 5 cm i.d.) was maintained at 50 °C. A magnetic stirrer maintained a consistent stir rate for all runs. The reactor was charged with 400 mL of an FeIIIHEDTA solution containing 1% by weight of iron at an initial pH of 7.5. Hydrogen sulfide was controlled using a mass flow meter at a rate equivalent to 400 [(mol of total Fe)/(mol of H2S)]/ min. Nitrogen was used to dilute the H2S stream just prior to entering the reactor solution through a sparger at the bottom of the reactor. Air (or oxygen) was sparged into the reactor through a separate line using a rotometer. The pH of the solution was maintained automatically by use of two syringe pumps, one containing sulfuric acid and the other 50% NaOH. Localized ferric hydroxide precipitation was minimized during the slow addition of the caustic solution by addition through narrow tubing into the well-stirred solution. Since the initial amount of ligand was not greatly reduced during the degradation, it is believed that any local Fe(OH)3 precipitate was quickly rechelated. Total iron concentration was checked at the end of the run to ensure that it remained constant. Total amounts of sulfuric acid and NaOH solutions never amounted to more than 1% of the initial reactor volume. Fe(II) was measured throughout the run by monitoring the absorbance of the complex formed by dilution in a solution of 1,10-phenanthroline. For each experiment, five cycles of ferric reduction/ferrous oxidation were run in which the percentage of Fe(II) cycled between 0 and 80% of the total iron. * To whom correspondence should be addressed. Tel.: 979-238-3240. Fax: 979-238-5183. E-mail:
[email protected].
Ligand analysis was performed by liquid chromatography. Regenerator Modeling. The regenerator modeling for a bubble column was reported previously.3 The (dynamic) mass balances for oxygen in the gas-phase reacting with ferrous chelate in the liquid phase are kla g ug ∂cAg ∂cAg ∂2cAg g ) DAX 2 mcA√1 + Ha2 ∂t ε ∂z ε ∂z g g
(1)
∂cB 4kla ul ∂2cB ∂cB l ) DAX mcg √1 + Ha2 2 ∂t (1 ε ) ∂z (1 - εg) A ∂z g (2) where A refers to the Fe(II) chelate and B to dioxygen. Other symbols are defined at the end of this paper. The above differential equations are numerically solved in a Pascal program. Inputs to that program are calculated using equations shown in ref 3. The program calculates the concentration of O2 in the gas phase as well as the concentration of Fe(II) in the liquid phase for any point in the reactor. One of the variables in the program is a time constant, which is increased until the outlet concentrations reach steady state. When steady-state performance has been assured, the output, which involves concentrations at various column heights, are exported to an Excel spreadsheet. The spreadsheet converts the gas-phase dioxygen concentrations (which are originally calculated at the average column pressure) to a local pO2 for each reactor segment (the reactor has been divided into 160 segments for the calculations). Average pO2 for each reactor configuration refers to the average pO2 for all the segments. Results and Discussion In practice, most iron chelate based desulfurization plants use a complex proprietary mixture of ligands, but the ligands HEDTA, EDTA, and NTA are most often used as models. This report concentrates on work with HEDTA and NTA. The first indication that the degradation of ligand might depend on O2 partial pressure came from the work of Wubs4 in which the stoichiometric factor, ν, for Fe(II) with dioxygen was seen to decrease from the theoretical value of 4.0 to values near
10.1021/ie900904s CCC: $40.75 2009 American Chemical Society Published on Web 10/05/2009
Ind. Eng. Chem. Res., Vol. 48, No. 23, 2009
Figure 1. Evidence for mass transfer limited ligand degradation in the oxidation of FeII HEDTA. r ) relative rate of air sparging; pO2 in atmospheres.
3.0. Such decreases were seen at higher pO2 values as well as at lower pH values (lower pH has been documented to cause increased ligand degradation).5 Such a discrepancy in ν was thought to occur because the dioxygen was used not only for Fe(II) to Fe(III) conversion but also for the oxidation of the ligand. Equations 3 and 4 show the overall reaction for the conversion of aminocarboxylates to the main products of ligand degradation: carbon dioxide and oxalate. 1.5O2 + H+ + R1R2NCH2COO- f R1R2NH + 2CO2 + H2O (3) O2 + R1R2NCH2COO- f R1R2NH + C2O42- + H+ (4) The following equations show the overall reaction for producing CO2 and oxalate from a hydroxyethyl group, such as that found on the ligand HEDTA. 2.5O2 + R1R2NCH2CH2OH f R1R2NH + 2CO2 + 2H2O (5)
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Fe L + O2 h [FeL-O2]
(7)
[FeL-O2] f FeIIIL + O2-
(8)
FeII|L + 2H+ + O2- f FeIIIL + H2O2
(9)
FeIIL + H2O2 f FeIIIL + HO• + OH-
(10)
FeIIL + HO• f FeIIIL + OH-
(11)
II
Other mechanisms are also possible involving a dinuclear oxygen complex and possibly metal hydroperoxide intermediates.5 In any case, there is only one reaction between Fe(II) and dioxygen with one of the subsequent reduction products (i.e., hydroxyl radical) believed to be most reactive toward the ligand. In this scenario, a “major” limitation on the availability of dioxygen to FeIIL should retard both degradation and regeneration to FeIIIL. However, a “minor” limitation of dioxygen could allow more regeneration to occur via reactions 9-11. This could be thought of as a “titration” of the Fe(II), allowing in only the dioxygen that is needed for regeneration. Dioxygen and the dioxygen reduction intermediates compete with each other for reaction with Fe(II). Starving the system of dioxygen will force the intermediates to react preferentially with Fe(II), rather than with the ligand. Fe(II) will be more prone to react with the intermediates rather than with a new dioxygen molecule (a process which would produce more dioxygen reduction intermediates). The high degree of oxidation of the ligand makes it most probable that another dioxygen molecule is involved in the reaction. If hydroxyl radical is involved the first reaction would probably produce a carbon centered ligand radical. LH + HO- f L• + H2O
(12)
The following scheme shows a possible mechanism involving additional dioxygen which would account for the presence of keto intermediates which have been detected in solution.
2O2 + R1R2NCH2CH2OH f R1R2NH + C4O42- + 2H+ + H2O (6) Several experimental runs were made in which the partial pressure and interfacial area of dioxygen in the regeneration gas were changed. All solutions started with the same concentration of ferrous HEDTA. Air was sparged until no more Fe(II) could be detected. The amount of degradation was measured by direct analysis of the ligand. By changing the partial pressure of O2 or the sparging gas rate, the rate of Fe(II) oxidation could be varied in a way that was predictable according to the regenerator model. Though increasing the sparging gas rate should increase the interfacial area, this correlation was not quantified. As partial pressure or gas rate was reduced, the Fe(II) oxidation rate decreased. Though at the end of the reaction the same number of moles of Fe(II) was oxidized, the amount of ligand degraded was substantially reduced as shown in Figure 1. It is most probable that the regeneration reaction must begin with some degree of complex formation between the ferrous chelate, FeIIL (L ) aminocarboxylate ligand) and dioxygen. This is shown below as a mononuclear (one iron center) complex, followed by reduction to form superoxide, hydrogen peroxide, and hydroxyl radical.
Regenerator Modeling Cases. Figure 1 showed that a regenerator designed to limit O2 dissolution and hence the Fe(II) oxidation rate (while allowing for the same degree of Fe(II) oxidation) could reduce associated ligand degradation. Most iron chelate based plants operate on air and do not have inert diluent gas available to reduce the partial pressure of O2. In addition,
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Figure 2. Fe(II) and pO2 profiles along reactor for case A. Table 1. Results of Regenerator Simulations case
ligand
reactor dimens (m)
A B
HEDTA HEDTA
C
HEDTA
D E
HEDTA HEDTA
F
NTA
4.93 × 1.078 1. 2.02 × 1.078 2. 2.02 × 1.078 1. 1.245 × 1.078 2. 1.245 × 1.078 3. 1.245 × 1.078 11.2 × 0.6819 1. 5.1 × 0.6819 2. 5.1 × 0.6819 6.35 × 1.078
a
ug (m/s)
ul (m/s)
av pO2 (atm)
deg coeff (g- · mol- · atm/m3)
3 3
0.0762 0.0762
0.00937 0.00937
0.1950 0.1968
0.2659 0.2120
3
0.0762
0.00937
0.1999
0.1990
1.2 1.2
0.0762 0.0762
0.0234 0.0234
0.1826 0.1848
0.2184 0.1967
3.0
0.0762
0.2107
0.2870
air stoichiometrya
Ratio of air used to stoichiometric requirement.
bubble column regenerator designs must maintain a minimum superficial velocity because of the need to keep elemental sulfur from falling to the bottom of the reactor. Use of full stream filtration (all sulfur removed prior to the regenerator) would allow for lower velocities to be used. The bubble column model was used to test potential effects of column dimensions on localized O2 concentrations and degradation. Case A: Base Design Case. The base case regenerator was sized for a plant with a capacity of one long ton per day of sulfur employing iron-HEDTA at a total iron concentration of 1.0 wt %. The specific gravity was assumed to be 1.20. The air flow rate was set at 0.25 ft/s (0.0762 m/s). The inlet concentration of Fe(II) was 0.7 wt %, and the desired outlet concentration was 0.3 wt %. The temperature used was 120 °F (322 K). Air was supplied at three times the stoichiometric amount. From the desired Fe(II) conversions, the air stoichiometric factor, and the linear air velocity, a cross-sectional area was calculated which gave a reactor diameter of 3.537 ft (1.078 m). From the diameter along with the plant’s sulfur capacity, the liquid velocity was calculated as 0.0308 ft/s (0.00937 m/s) The regenerator model was run using various reactor heights until the desired outlet concentration of Fe(II) was achieved. Figure 2 shows the concentration profiles for dioxygen and Fe(II) at various column heights. Other simulations are summarized in Table 1. In cases B and C, one large reactor was replaced with two and three reactors, respectively (same diameter as the reactor in case A). The
objective here was to avoid high inlet pressures caused by the large hydrostatic head of the column. This was achieved by using multiple short columns (separate pumps in-between). Though the inlet pressure followed the order A > B > C, the average pressure over the reactor(s) was approximately the same for each case, due to differences in reaction rates and column back-mixing. Different reactor configurations were accompanied by different slopes and curve shapes for Fe(II) reactor profiles. A greater slope at any point in the reactor means that more oxidation is occurring in that section of the reaction. Calculations were made to compare the slope in each section to the sum of the slopes for all sections. This made it possible to determine the relative amount of Fe(II) oxidation occurring within that section. Since the degree of expected degradation should depend on the amount of Fe(II) oxidized and is believed to be first order in pO2, the relative amount of degradation in each section was calculated by multiplying the change in Fe(II) concentration within that section by the average pO2 in that section. The resulting number is referred to as the degradation coefficient n j 2)i, ∆[FeII]i(pO and is defined as follows: coefficient ) ∑i)1 j where n ) number of segments in the reactor, (pO2)i ) average pressure in each segment, and ∆[FeII]i ) change in ferrous concentration in each segment. Coefficients were totaled for simulations of multiple reactor configurations. Such analysis shows that configuration B should have about 20% less degradation than for base case A.
Ind. Eng. Chem. Res., Vol. 48, No. 23, 2009
Utilization of three reactors (case C) results in only a small improvement over two reactors. In cases D and E, the stoichiometric amount of air has been decreased to 1.2 (20% excess of what is required) by keeping the superficial velocity the same as the base case (this requires a narrower reactor). Case D shows that, for a single reactor configuration, reduction of excess air (maintaining superficial velocity) should result in about a 20% reduction in ligand degradation. As in cases A-C there is also additional benefit in going to multiple reactors (case E). For case F, nitrilotriacetic acid (NTA) was used as the ligand. The value for the rate constant for the reaction of FeIIL and O2 was reduced from 0.0838 (for HEDTA) to 0.0401 m6 mol-2 s-1. This required a taller column. In spite of cutting the reaction rate constant in half, the reactor for case F was only 29% taller than case A. This is because the rate of regeneration is limited by the O2 mass-transfer rate as well as by the reaction rate. Case F also resulted in a higher degradation coefficient, though these coefficients should not be used to compare different ligands. The model cases reported here did not consider complete Fe(II) regeneration because it has been found that maintaining higher Fe(II) concentrations during the regeneration results in reduced degradation.6 Since control of Fe(II) is accomplished by reducing the amount of air to the regenerator (or reducing residence time), it could be argued that Fe(II) control works by implementing a mass-transfer limitation to the degradation reaction. It has long been thought that Fe(II) control functions by radical scavenging. Hydroxyl radical scavengers have been found to suppress the ligand degradation reactions.7 Fe(H2O)62+ and ferrous chelates of nitrilotriacetic acid and ethylenediaminetetraacetic acid all exhibit high reactivities toward hydroxyl radical.8 Analysis of previous FeIIL oxidation rate studies4,9 has shown that mass-transfer control of the reaction can switch to kinetic control if the ferrous concentration is lowered or if the liquid film mass-transfer coefficient is increased. Thus, the effects demonstrated and predicted in this work may be valid only under specific reaction conditions. Recent developments in flue gas CO2 capture have suggested that amine degradation reactions in that application may be limited by the mass transfer of dioxygen into solution.10 Since transition metal catalysis of this autoxidation may involve formation of dioxygen reduction intermediates,11 part of this mass-transfer limitation may involve mechanisms similar to those proposed here. Conclusions A model for bubble column regeneration for an iron chelate dehydrosulfurization process was used to predict local partial pressures of dioxygen for a number of reactor configurations. From these local partial pressures and the amount of Fe(II) being oxidized in each reactor segment, relative amounts of ligand degradation were estimated. Such analysis predicts several changes in regenerator design which could substantially reduce the amount of chemical usage in a chelated iron plant. Changes include reduction of the stoichiometric excess of air, use of multiple reactors, and recycling of spent air. Symbols a ) specific area (1/m) c ) concentration (mol/m3)
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DA/B ) diffusion coefficient (m /s) Dax ) dispersion coefficient (m2/s) EA ) enhancement factor EA∞ ) instantaneous enhancement factor Ha ) Hatta number He ) Henry coefficient (Pa · m3/mol) kl ) mass-transfer coefficient (m/s) k ) reaction rate constant (varies) m ) distribution coefficient p ) pressure (Pa) t ) time (s) T ) temperature (K) u ) superficial velocity (m/s) y ) mole fraction z ) distance (m) 2
Subscripts/Superscripts A or O2 ) oxygen B ) ferrous chelate g ) gas phase l ) liquid phase ax ) axial i ) interface Greek Letter ε ) hold-up
Literature Cited (1) Kohl, A.; Nielsen, R. Gas Purification, 5th ed.; Gulf: Houston, TX, 1997. (2) Chen, D.; Motekaitis, R. J.; Martell, A. E.; McManus, D. Oxidation of H2S to S by Air with Fe(III)NTA as a Catalyst: Catalyst Degradation. Can. J. Chem. 1993, 71, 1524. (3) Jongsma, F. J.; Demmink, J. F.; Bedell, S. A.; Beenackers, A. A. C. M. Regeneration of Iron Chelate Solutions in a Multistaged Bubble Column ReactorsComparison of Single and Multistaged Bubble Columns, Preprint Paper, 8th Gas Research Institute Sulfur Recovery Conference, Oct. 12-15, 1997. (4) 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. (5) Liu, X.; Sawyer, D. T.; Worley, C. M.; Bedell, S. A. Ligand Degradation in the Iron/Dioxygen-Induced Dehydrogenation of H2S, Preprint Paper, 7th Gas Research Institute Sulfur Recovery Conference, Sep. 24-27, 1997. (6) Jeffrey, G. C.; Myers, J. D. Process and Composition for the RemoVal of Hydrogen Sulfide from Gaseous Streams, U.S. Patent 4774071, Sep. 27, 1988. (7) Bedell, S. A. Stabilized Chelating Agents for RemoVing Hydrogen Sulfide, U.S. Patent 4891205, Jan. 2, 1990. (8) Neta, P.; Huie, R. E. Rate Constants for Reactions of Peroxyl Radicals in Fluid Solutions. J. Phys. Chem. Ref. Data 1990, 19, 413. (9) Demmink, J. 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, 1989. (10) Goff, G. S.; Rochelle, G. T. Monoethanolamine Degradation: O2 Mass Transfer Effects under CO2 Capture Conditions. Ind. Eng. Chem. Res. 2004, 42, 6400. (11) Bedell, S. A. OxidatiVe Degradation Mechanisms for Amines in Flue Gas Capture Lessons Learned from Other Gas Treating Processes, Preprint Paper, 9th Greenhouse Gas Control Technologies Conference, Washington, D.C., November 2009.
ReceiVed for reView June 2, 2009 ReVised manuscript receiVed September 21, 2009 Accepted September 21, 2009 IE900904S
