Reaction with iron, kinetics

Iron frequently has been postulated to be an important electron acceptor for oxidation of sulfide (58, 84,119, 142, 152). Experimental and theoretical studies have demonstrated that Fe(III) will oxidize pyrite (153-157). Reductive dissolution of iron oxides by sulfide also is well documented. Progressive depletion of iron oxides often is coincident with increases in iron sulfides in marine sediments (94, 158, 159). Low concentrations of sulfide even in zones of rapid sulfide formation were attributed to reactions with iron oxides (94). Pyzik and Sommer (160) and Rickard (161) studied the kinetics of goethite reduction by sulfide thiosulfate and elemental S were the oxidized S species identified. Recent investigations of reductive dissolution of hematite and lepidocrocite found polysulfides, thiosulfate, sulfite, and sulfate as end products (162, 163). 

Aquocopper(I) complexes are fairly powerful reducing agents and the kinetics of their reactions with iron(III) (275), vanadium(IV) (312), cobalt(III), and mercury(II) (113) have been studied. Further, the role of copper(I) species in the copper(II)-catalyzed reduction of cobalt(III) by vanadium(II) (112) has been confirmed with the reduction of... 

The oxidation of [ Fe(CO)2Cp 2] is irreversible and the potential is very electrode dependent (e.g., 1/2 = 0.19 V at carbon, 1.0 V at platinum, in MeCN). CPE (336) or reaction with iron(III) salts (337) in the presence of a ligand L provides a good route to [Fe(CO)2LCp]+ (L = acetone, NCMe, P donor, etc.). Coulometry reveals that the oxidative cleavage reaction requires the loss of two electrons per dimer (336). However, kinetic studies of the reaction with [RuCl2(bipy)2]+ or [ Fe(CO)Cp 4] + show the ratedetermining step to involve one-electron transfer (338). 

Another way to slow substitution is to covalently bond a ligand to silica— compare bonding ligands to monoclonal antibodies above. Oxine bound to silica reacts much more slowly with Al, for example, than when it is in solution. Another example of slow substitution at Al is related to the indium chemistry mentioned above, involving its slow transfer from its transferrin complex by reaction with iron(III)-citrate to form the much more stable combination of iron-transferrin and aluminum-citrate complexes. Further examples of slow substitution kinetics involving ferritin will be found in the iron(III) section (Section 8.3.4). 

RIchtsmeler S C, Parks E K, Liu K, Polo L G and Riley S J 1985 Gas phase reaction of Iron clusters with hydrogen. I. Kinetics J. Chem. Rhys. 82 3659... 

The oxidation reaction (eq. 2) is an irreversible reaction with first order in oxygen and smmd ordo in iron chelate. The kinetic constants at various temperatures were determined using a penetration theory based expression derived by De Coursey [4] and the results are presented in Fig. 2. 

The last reaction is the most favored of these three. The actual occurrence of the reactions with elemental phosphorus or phosphorous trichloride as products has been explained to be due to kinetic reasons. The thorium present in the ore volatilizes in the form of thorium tetrachloride (ThCl4) vapor other metallic impurities such as iron, chromium, aluminum, and titanium also form chlorides and vaporize. The product obtained after chlorination at 900 °C is virtually free from thorium chloride and phosphorous compounds, and also from the metals iron, aluminum, chromium, and titanium. 

Lewis Bases. A variety of other ligands have been studied, but with only a few of the transition metals. There is still a lot of room for scoping work in this direction. Other reactant systems reported are ammoni a(2e), methanol (3h), and hydrogen sulfide(3b) with iron, and benzene with tungsten (Tf) and plati num(3a). In a qualitative sense all of these reactions appear to occur at, or near gas kinetic rates without distinct size selectivity. The ammonia chemisorbs on each collision with no size selective behavior. These complexes have lower ionization potential indicative of the donor type ligands. Saturation studies have indicated a variety of absorption sites on a single size cluster(51). 

Kinetics and activation parameters for NO reactions with a series of iron(II) aminocarboxylato complexes have been obtained (Table II) in aqueous solution (31). Rate constants for these reactions ranged from 105 to 108M-1s-1 for the series of iron(II) complexes studied. The reactions of NO with Fen(edta) (edta = ethylenediaminetetraacetate) and Fen(Hedtra) (Hedtra = hydroxyethylenediaminetriacetate) yielded activation volumes of +4.1 and +2.8 cm3 mol-1, respectively and were assigned to a dissociative interchange (Id) mechanism (31b). All of the iron(II) aminocarboxylato complexes studied followed a similar pattern with the exception of the Fen(nta) (Nta = nitriloacetic acid) complex which gave a AV value of —1.5 cm3 mol-1. The reaction of this complex with... 

The NO/NO+ and NO/NO- self-exchange rates are quite slow (42). Therefore, the kinetics of nitric oxide electron transfer reactions are strongly affected by transition metal complexes, particularly by those that are labile and redox active which can serve to promote these reactions. Although iron is the most important metal target for nitric oxide in mammalian biology, other metal centers might also react with NO. For example, both cobalt (in the form of cobalamin) (43,44) and copper (in the form of different types of copper proteins) (45) have been identified as potential NO targets. In addition, a substantial fraction of the bacterial nitrite reductases (which catalyze reduction of NO2 to NO) are copper enzymes (46). The interactions of NO with such metal centers continue to be rich for further exploration. 

Iron(III)-catalyzed autoxidation of ascorbic acid has received considerably less attention than the comparable reactions with copper species. Anaerobic studies confirmed that Fe(III) can easily oxidize ascorbic acid to dehydroascorbic acid. Xu and Jordan reported two-stage kinetics for this system in the presence of an excess of the metal ion, and suggested the fast formation of iron(III) ascorbate complexes which undergo reversible electron transfer steps (21). However, Bansch and coworkers did not find spectral evidence for the formation of ascorbate complexes in excess ascorbic acid (22). On the basis of a combined pH, temperature and pressure dependence study these authors confirmed that the oxidation by Fe(H20)g+ proceeds via an outer-sphere mechanism, while the reaction with Fe(H20)50H2+ is substitution-controlled and follows an inner-sphere electron transfer path. To some extent, these results may contradict with the model proposed by Taqui Khan and Martell (6), because the oxidation by the metal ion may take place before the ternary oxygen complex is actually formed in Eq. (17). 

The three rate constants for Eq. (98) correspond to the acid-catalyzed, the acid-independent and the hydrolytic paths of the dimer-monomer equilibrium, respectively, and were evaluated independently (107). The results clearly demonstrate that the complexity of the kinetic processes is due to the interplay of the hydrolytic and the complex-formation steps and is not a consequence of electron transfer reactions. In fact, the first-order decomposition of the FeS03 complex is the only redox step which contributes to the overall kinetic profiles, because subsequent reactions with the sulfite ion radical and other intermediates are considerably faster. The presence of dioxygen did not affect the kinetic traces when a large excess of the metal ion is present, confirming that either the formation of the SO5 radical (Eq. (91)) is suppressed by reaction (101), or the reactions of Fe(II) with SO and HSO5 are preferred over those of HSO3 as was predicted by Warneck and Ziajka (86). Recently, first-order formation of iron(II) was confirmed in this system (108), which supports the first possibility cited, though the other alternative can also be feasible under certain circumstances. 

It is ironic that the protonic acids, which were chosen originally by Pepper as apparently offering the best prospect of simplest kinetics, have turned out to be fundamentally unsuitable for kinetic work because of the multiplicity of their possible reactions with monomers, and the formation of conjugate anions. 

Use of the kinetic advantage method thus points clearly to the occurrence of chemical catalysis with the low-valent metalloporphyrins. This is confirmed by repeating, with iron(I) octaethylporphyrin and cobalt (I) etioporphyrin, the stereochemical experiments carried out earlier with the anion radical of 1,4-diacetylbenzene. Complete stereospecificity is observed in both cases The meso isomer of 4,5-dibromooctane is converted totally into the c/.v-olcfin the d,l isomer is converted totally into the trans-olefin. The reaction again exhibits a clear antiperiplanar preference. 

The obtained obs values are reported as a fimction of the complex concentration (Fig. 8), and a good linear correlation between obs and the complex concentration was observed for both oxidation forms of iron complexes. From the slope of the plot of obs vs. catalyst concentration the catalytic rate constants ( cat) i29) were determined to be (3.7 0.5) X 10 M- s and (3.9 0.5) x 10 M s for [Fe (dapsox)(-H20)2]C104 and [Fe (H2dapsox)(H20)2](NOs)2, respectively 49). It is important to note that, it does not matter whether we start from the Fe(III) or Fe(II) form of the complex, identical spectral changes (Fig. 6a and 6b) and kinetic behavior (Fig. 8) for these two complexes is observed upon reaction with, which is consistent with the redox cychng of the complex during O2 decomposition (Scheme 9). 

Among comparative kinetic studies, the kinetic advantage method has been used systematically in several cases. It has been developed for the first time for investigating the ET versus 8 2 problem in the reaction of iron(i) and cobalt(ii) porphyrins with primary butyl halides (Lexa et al., 1981), yielding the corresponding a-butyl-iron(iii) and cobalt(iii) complexes according to the overall reaction (142). 

Fig. 16.10 Plot showing kinetics of C CljNOj reduction (fiUed circles) occurring in conjunction with increasing photon correlation spectrometry (PCS) count rates (open circles), which are indicative of particle formation, in reaction with O.SOmM Fe(ll) (pH 7.0). (For clarity, the symbols showing measured values of [C CljNOJ are connected point to point.) The other open symbols show PCS count rates in nonreaction mixtures (i.e., without C Cl NO ) containing either O.SOmM Fe(II) (pH 7.0) or O.SOmM Ca(ll) (pH 7.0). Reprinted with permission from Klupinski TP, Chin YP, Traina SJ (2004) Abiotic degradation of pentachloronitrobenzene by Fe(ll) Reactions on goethite and iron oxide nanoparticles. Environ Sci Technol 3S 4353-4360. Copyright 2004 American Chemical Society...

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