Oxidation kinetics of Iron(II) complexes of trans-1,2-diaminocyclohexanetetraacetate (cdta) with dissolved oxygen: Reaction mechanism, parameters of activation and kinetic salt effects
Introduction
The oxidation of iron(II) polyaminocarboxylate complexes with molecular oxygen (O2) and its derivatives (i.e., superoxides ), hydrogen peroxide (H2O2), hydroxyl radical ) has been studied for some time now in areas like biochemistry (i.e., superoxide dismutation, H2O2 decomposition—Lati and Meyerstein, 1978, Cohen et al., 1981, Bull et al., 1983, Rush and Koppenol, 1988, Rahhal and Richter, 1988, Seibig and van Eldik, 1997, Seibig and van Eldik, 1999), environmental science (i.e., catalytic removal of and , gas desulphurization—Brown and Mazzarella, 1987, Sada et al., 1987, Zang and van Eldik, 1990, Wubs and Beenackers, 1993, Demmink and Beenackers, 1997, Sharma et al., 2004) and water science (Liang et al., 1993). Recently, gas desulphurization processes based on ferric chelate chemistry have received increasing attention in the area of hydrocarbon processing (i.e., LO-CAT process: U.S. Filter Engineered Systems, Hardison and Ramshaw, 1992; Sulfint process, Mackinger et al., 1982). Application of such methods for hydrogen sulfide removal in the pulp and paper industry would differ on the flue gas composition because of the large amounts of oxygen associated with the atmospheric emissions. In line with this latter category of atmospheric emissions, Iliuta and Larachi (2003) recently established a modelling framework characterizing the potential of a bifunctional redox process where an iron(III) chelate ( where L denotes an organic ligand of charge) is used to throttle hydrogen sulfide emissions (Eq. (1)) whilst dissolved oxygen regenerates simultaneously the iron(II) chelate product into the active iron(III) form (Eq. (2)). Emphasizing on the regeneration step of such process, this work will strictly focus on the kinetics and mechanism of iron(II) chelate oxidation with dissolved oxygen as the source oxidant.
Experimental efforts in the literature regarding the reaction kinetics and procedures are largely summarized in Wubs and Beenackers (1993) work. Sada et al. (1987) proposed a reaction scheme based on experiments conducted in a bubble column using iron(II) nitrilotriacetate (nta) and iron(II) ethylenediaminetetraacetate (edta) complexes. Kinetic measurements were acquired based on O2 gas/liquid mass transfer. They established the reaction rate to be first-order with respect to O2 while ranging between 0 and 1 for . However, the proposed rate expression becomes irrelevant for general situations since the O2 concentration was regarded to be negligible in contrast to the concentration. In addition, this assumption was not well established and could have led to misinterpretation considering that other studies on the subject report a reaction order between 1 and 2. Brown and Mazzarella (1987) developed a reaction scheme involving 4 forward and 1 reverse steps (Eqs. (3)–(6)).This was obtained from experiments conducted in O2-saturated solutions of iron(II) aminopolycarboxylate complexes, in which a rotating pyrolitic graphite electrode was immersed to monitor the oxidation rate.
The final steps (Eqs. (5)–(6)) were considered very fast due to stronger oxidants and so the rate law turned out to be first-order with respect to O2 while evolving from first to second order with respect to . The latter partial order would indeed depend on the concentration and the relative strength of the reverse step with the subsequent forward step (, Eq. (4)). Later on, Zang and van Eldik (1990) proposed a slightly different reaction mechanism hypothesizing that and binds together to form a non-oxidized intermediate which is consumed via three parallel steps (Eqs. (7)–(9)). Superoxides ( or ) and would then oxidize molecules following reaction steps similar to those depicted by Eqs. (4)–(6).
Although molecular oxygen is a weak oxidant, it seems unusual that an intermediate like would accumulate to such level that Reaction (9) could take place significantly. Unless large amounts of are used, should be considered an activated complex immediately forming or . Still, Wubs and Beenackers (1993) came with the same reaction scheme based on experiments conducted in a flat-interface stirred cell reactor containing iron(II)–edta or iron(II) (hydroxyethyl)ethylenediaminetetraacetate (hedta) solutions. Demmink and Beenackers (1997) studied the reaction with iron(II)–nta solutions. Kinetic measurements were obtained by a manometric method monitoring the head space pressure in the stirred cell. They noted a slight but consistent increase in consumption from expected stoichiometry. Accordingly, they modelled an iron chelate degradation step into their overall scheme stipulating that a fraction of the hydrogen peroxide produced during the reaction would attack functional groups (i.e., ethylene moiety of edta, carboxylate groups) splitting the iron from the organic matrix. This phenomenon was previously ascribed to the formation of hydroxyl free radicals (Chen et al., 1995).
More recently, Seibig and van Eldik, 1997, Seibig and van Eldik, 1999 revisited the reaction with help of stopped-flow units connected to a diode array spectrophotometer. Their study has put a lot of importance on iron(II) chelate complex chemistry making clear that different species will result in different reactivity with . For example, it was alleged that iron(II)–edta having one protonated acetate arm self-arranges into a pentagonal bipyramidal structure, which is believed to bind much faster than mono-capped trigonal-prismatic species such as nonprotonated and diprotonated iron(II)–edta. Accordingly, equilibrium reactions illustrating the pH-dependence of different complexes were added to the reaction scheme without significantly modifying the elementary steps from the ones proposed by Zang and van Eldik (1990).
Overall assessment of previous studies leads to the following conclusions with respect to pH. The reaction rate is normally pH-independent in the alkaline-to-neutral region (DeBerry, 1997). A reaction enhancement is observed at lower pH and a maximum is reached in the vicinity of depending on the ligand in use (Sada et al., 1987, Zang and van Eldik, 1990, Seibig and van Eldik, 1997, Seibig and van Eldik, 1999). This is likely ascribable to the formation of more reactive complexes as mentioned by Seibig and van Eldik, 1997, Seibig and van Eldik, 1999. The oxidation rate then falls drastically at pH lower than 2 where dechelation is triggered. The exact pH where the iron(II) aminopolycarboxylate complex will start to split into its building species (i.e., and ) is characterized by the formation constant (Smith et al., 1985a, Smith et al., 1985b; Smith and Martell, 1987) along with the ligand deprotonation constants. The lack of reactivity at very low pH supports the fact that hydrated iron(II) ions weakly react with .
Up till now, all known experimental setups (i.e., bubble column, stirred cell) were used to investigate the oxidation rate using spectrophotometrical, electrochemical or pressure methods. In no case was the dissolved oxygen concentration monitored during the reaction apart from being known at the reaction start or maintained at saturation using vigorous sparging. In this work, an experimental setup was developed to quantify , and simultaneously in a completely filled stirred cell reactor using trans-1,2-diaminocyclohexanetetraacetatic acid (cdta) as the ligand. Concentration measurements from a UV–Vis spectrophotometer (iron(II)–cdta, iron(III)–cdta) and a polarographic electrode were used to clarify the reaction mechanism and determine thermodynamic properties (parameters of activation, kinetic salt effects) with help of elements discussed in previous studies. The experimental settings involve broad intervals of temperatures , ionic strengths and pH . From an engineering perspective, this work focuses on the importance of knowing the effect of every possible aqueous setting on the oxidation rate in order to prepare accurate design procedures for scrubbing processes exploiting the iron chelate chemistry in the mitigation of the odors associated with the emissions of total reduced sulphurs in the pulp and paper atmospheric air emissions.
Section snippets
Experimental section
Kinetic experiments were conducted in a hybrid aluminum/Plexiglas stirred cell of 12.7 cm i.d. and 25.4 cm height (Fig. 1) as described in a previous work (Piché and Larachi, 2005) allowing some modifications: (1) two extra six-blade turbine stirrers (6.35 cm i.d.) for a total of four were added in the upper section to maximize liquid mixing in the stirred cell reactor; (2) a Clark-type polarographic electrode (Omega DOB-930) was placed in the cell with its tip positioned at about 7 cm
O2, iron(II)–cdta and iron(III)–cdta quantification
Dissolved oxygen saturation depends foremost of the surrounding pressure, temperature and ionic composition. The operating pressure was never recorded but was assumed to be near atmospheric. A thermocouple ring attached onto the polarographic electrode allowed automatic temperature correction. In contrast, no conductivity correction was directly imposed from the meter. Instead, salting-out constants (; , , —Lang and Zander, 1986)
Kinetics in alkaline solutions
Forty-three kinetic trials were used as basis for developing a consistent reaction mechanism for operation in alkaline solutions (–12). Temperature and ionic strength were maintained at and of NaCl per , respectively. In these conditions, only one iron(II)–cdta complex forms (Seibig and van Eldik, 1999) while the and products co-exist in equilibrium. Initial and concentrations were kept in
Conclusion
Oxidation of iron(II) trans-1,2,-diaminocyclohexanetetraacetate (cdta) complexes with molecular oxygen as the source oxidant was studied in alkaline (mechanism and kinetics, activation parameters, kinetics salt effects) and acidic (kinetics) media. and are the only iron(II)–cdta and superoxide species formed at pH larger than 8. So it led to the elaboration of a reaction rate model (Eq. (22)) coupled with two apparent rate constant functions (Eqs. (34)–(35)) characterized by
Notation
iron concentration, oxygen solubility, salt concentration, ionic strength, limiting step apparent rate constant in Eq. (22), salting out constant, equilibrium constant of activated complex, Eq. (16) n complex coordination r rate constant, S stoichiometric ratio t time, s T absolute temperature, K Greek letter activity coefficient Sub/superscripts 0 initial #1 first activated complex
Acknowledgement
Financial support from the Natural Sciences and Engineering Research Council of Canada (NSERC) and the Fonds Québécois de la Recherche sur la Nature et les Technologies is gratefully acknowledged.
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