J. Chem. SOC., Faraday Trans. I, 1986,82, 2345-2352 Dissolution of Cobalt Ferrites by Thioglycolic Acid Miguel A. Blesa," Albert0 J. G. Maroto and Pedro J. Morando Departamento Quimica de Reactores, Comision Nacional de Energia A tbrnica, Avenida del Lihertador 8250, 1429 Buenos Aires, Argentina The dissolution of cobalt ferrites CoZFe3-Z04 by thioglycolic acid inxiolves the chemisorption of thioglycolate anion onto Fe"I ions of the solid, followed by an electron transfer from the ligand to the metal ion and subsequent release of FeII. Kinetic data suggest that two adjacent Fe"'-L sites evolve to two FelI+L,. Substitution of Co" for Fe" does not bring about any noticeable change in the kinetics for x < 0.6. For larger values of x, the early mechanism of dissolution changes, suggesting that electron hopping within the octahedral sites may produce a chain dissolution of FeI' for each single original electron transfer from thioglycolate.Rate data in the presence of exogenous FeI' are also discussed. ~~~ The affinity of the thioglycolate anion towards FelI and FeIII is well known. In the former case the complexes have been characterized structurally and thermodynamically [ref. ( 1) and references therein] ; Felll-thioglycolate complexes, on the other hand, are unstable, and decompose to FeII and disulphur compounds.i The high stability of the complexes, as well as the deep colour characteristic of some of them, was the basis of the use of thioglycolic acid in the determination of iron in ores and insoluble oxides.2 Implicit in this use was the high dissolving potency of thioglycolic acid for iron1I1 o x i d e ~ .~ In a former paper4 we reported on the influence of solution variables (pH and acid concentration) on the dissolution of magnetite in thioglycolic acid solutions, thus putting forward evidence for the role of surface Fe-thioglycolate species in the dissolution mechanism. Now we have explored the characteristics of the dissolution process in the cobalt ferrite-thioglycolic acid system. By using ferrites of formula Cox Fe3-x04 (0 < x < 0.9) we have now been able to put forward evidence for some subtle characteristics of the dissolution process that are dependent on the structure of the solid. Since our original paper on Fe,O,-thi~glycolate,~ we have shown several instances of the 'reductive dissolution' mechanism of magnetite, mainly by FelI cornplexe~~-~~ and similar reductive mechanisms have been investigated and clarified through the work of Sellers and c o ~ o r k e r s .~ ~ - ~ ~ The mechanism involves an interfacial electron transfer from an adsorbed species to an FerIr surface state as the crucial point in the mechanism.lo Elegant electron microscopy by Sellers and coworkersi5 has shown that the FeI1 surface states thus formed are highly reactive and dissolve readily, thus justifying the assumption of rate control by the electron-transfer process. The abovementioned work has not put forward evidence of any kinetic manifestation of the semiconducting properties of magnetite. The dissolution process is akin to phenomena such as photocurrents at the semiconductor-electrolyte solution interfacei83 l9 that can be understood only on the basis of an adequate description of the electronic surface structure of the semiconductor.Segall and coworkersz0 have advanced evidence of the importance of the semiconducting characteristics of NiO in its oxidative dissolution mechanism. Similar considerations were put forward by NiiZ1 and DiggleZ2 These authors have shown that the availability of minority or majority carriers at the interface may control the dissolution rate. It is well know that both the crystal and electronic structures of magnetite can suffer 78 2345 FAR 12346 Dissolution of Cobalt Ferrites important changes upon substitution. Thus, magnetite presents an inverse spinel ~tructure,,~ with FelI1 ions distributed evenly in tetrahedral and octahedral sites and FerT located in octahedral sites only.There is a fast electron-hopping between [FeIII], and [FeTT], that is responsible for the high electrical conductivity of magnetite.24 Upon substitution of ColI for Feil, cobalt distributes unevenly in octahedral and tetrahedral sitesz5 and when x reaches ca. 0.7 the electrical conductivity drops to very low values26 because there are no close-neighbour pairs [Fe1i1-Fe11]o.27 Our previous results had given some clues in the sense that electrical conductivity might have some non-trivial consequences in the dissolution p r o ~ e s s . ~ We now report data on the dissolution of cobalt ferrites that prove the involvement of this property in the dissolution mechanism.Experimental All reagents employed were of analytical purity and were used as provided. Ferrites were prepared as described for magnetite,2s but partially replacing Coil for Fe**. Mixtures of FeCl, * 4H,O and CoCl, * 6H,O were treated in boiling aqueous solution with ammonia in the presence of hydrazine and NaNO,; the resulting slurry was aged under continuous stirring for 30 min. The solid thus formed was rinsed several times with water, filtered off and dried in a desiccator at room temperature. The samples were characterized by chemical analysis, X-ray diffraction, scanning electron microscopy and specific surface area measurements. The solids were in all cases composed of cubic particles of average edge 0.1 ,urn and with a rather small extent of polydispersion.The specific surface area values ranged between 7 and 9 m2 g-l, in good agreement with the calculated geometric area for the particles. Kinetics experiments were performed as described previ~usly.~ The ferrite (usually 40 mg) was suspended in doubly distilled water in a magnetically stirred cylindrical beaker provided with a water jacket. The reaction was started by adding solutions of thioglycolic acid and sodium hydroxide (to obtain the desired pH); the volume of the reacting slurry was 170 cm3. Samples were taken periodically and poured into a large volume of water containing thioglycolic acid (TGA) and excess ammonia. This solution was filtered through a Nuclepore membrane (pore size 0.45pm) and the absorbance at 530 nm was measured in a Shimadzu UV-210A spectrophotometer.The amount of dissolved iron was then calculated from the calibration curve. Experiments carried out at various stirring rates were performed to ensure that the reaction was not under diffusional control. Results and Discussion The Shape of a vs. t Plots Plots of the fraction of ferrite dissolved as a function of time are shown in fig. 1 for various x values. In the range 0 6 x 6.0.6 neither the shape nor the time scale is noticeably dependent on x. On further increasing x, an induction period ensues. The change in the shape is not simply of the type that can be accounted for by shifting the time scale [e.g. plotting a =flt/lo.5)]zg and suggests, therefore, a change in the reaction mechanism, i.e. in the relative importance of the nucleation and growth stages.3o Our previous results on magnetite dissolution by TGA were interpreted contracting-sphere kinetics, i.e.assuming that nucleation was fast over all the solid s ~ r f a c e . ~ The cubic-root law holds well for cobalt ferrites with x < 0.6: see fig. 2. Except when noted, all quoted k values refer to the expression 1 -(1 -a)$ = k t . (1) Almost identical rate constants were obtained for x < 0.6: k = (8.2f0.4) x lop2 s-l.M. A . Blesa, A . J , G. Maroto and P. J . Morando 2347 1 .o a 0 . 5 0 5 10 15 20 25 tlmin Fig. 1. Dissolution fraction a us time profiles for various Co,Fe,-,O, at 70 "C and total thioglycolic acidconcentrationC,,, = 6.5 x mol dm-"pH 3.70; .,x = 0.16; 0 , x = 0.25; 0 , x = 0.50; B, x = 0.69. 0.6 m - h a I 0.4 & v I - 0 .2 0 A / i / 2 4 6 8 1 0 r/min Fig. 2. Linear plot of 1 -(1 -a); as function of time at 70 "C; A, pH 3.90; e, pH 3.32; C,,, = 6.5 x mol drn-,. Electron Microscopic Characterization of the Dissolution Process SEM was used in an attempt to distinguish between massive and localized attack on the surface of ferrite crystals. Massive dissolution is apparent from the photographs, however, because of the small size of the crystal and the rather low magnification available (30000), pitting cannot be ruled out. 78-22348 Dissolution of Cobalt Ferrites 7.0 d I N I WY 4: Y 5 -0 3.0 1 I I 3.0 1.0 5.0 PH Fig. 3. Influence of pH on the specific rate constant k ; C,,, = 6.5 x mol d ~ n - ~ ; t = 70 “C; 0, experimental; (-a -) calculated assuming first order on surface complex concentration, k = kE”,”[C, FeTGA)/ qwaFeTGA)] ; (---) calculated assuming second-order dependence on surface complex concentration, k = kg;[Cq_ FrTGA)/.q-a$&GA)].Surface complex concentration ratios were calculated according to ref. (4). Superscript ‘max’ in the above relations refers to maximum rate, i.e. pH 3.8. The Influence of Solution Variables on the Rate of Dissolution pH influences the rate in the same way as described before for magnetite.4 Fig. 3 shows the values of the dissolution rate constant ( k ) us. pH at 70 “C. The interpretation is as follows in a modified version of the reasoning given in ref. (4). The basic idea is that the data in fig. 3 simply describe the pH dependence of the adsorption equilibrium, e.g. the change in the concentration of surface complexes FeIII-TGA with pH.or its equivalent, &ds, calculated from Langmuir-type expression^.^ 9 31-33 AGids has been shown to exhibit maxima (or humps) at pH values close to the pK of the conjugate acid of the complexing anion.7* 34-36 In fact, the shape of the curve can be very asymmetrical, showing a not very pronounced decrease in the acidic branch and this can be modelled by assuming that the complexing anion replaces both OH- ions and H,O molecules bound to metal ions located in kinks of the solid surface :37 Adsorption of complexing anions onto metal oxides can be characterized by M(OH),-r(H,O) + A- + (1 - 5)H’ + MAC- + H 2 0 . (2) In eqn (2), 5 is the fraction of exposed M sites accounted for by H,O ligands.For high coverages, 5 -+ 1, a very shallow maximum and a very asymmetrical curve are expected, as actually found in several instance^.^^^^^^^ Thus, the pH dependence of the rate (fig. 3) is controlled by the dependence of the concentration of surface complexes on pH. In our former paper on the dissolution of magnetite,4 the pH dependence of adsorption was assumed according to a very simple model that predicts a symmetric curve, the width of which was determined by the acidity constants of the surface and of thioglycolic acid. Even on this assumption, a much better fit is obtained if the rate is assumed to be second-order with respect to surface complex concentration R = khom[ - Fe1I1-TGAI2. (3)M. A . Blesa, A . J . G. Maroto and P. J . Morando 2349 This is shown in the calculated profiles included in fig.3. Following the procedure outlined in ref. (4), a term of the form: was assumed to account for the pH dependence of surface Fe-TGA complexes; Ka, and Ki are the first acidity constants of thioglycolic acid and of Fe-OH: surface sites, respectively. The outermost profile in fig. 3 represents precisely the behaviour of the above expression, whilst the inner profiles represents the behaviour of the squared expression. In both cases the height of the maximum was chosen to fit the kinetic data. This result agrees well with mechanistic information from homogeneous aqueous systems; depending upon the experimental conditions, the rate of the internal redox reaction in Fe'II-thioglycolate complexes has been found to be or ~ e c o n d - o r d e r ~ ~ - ~ ~ in the complex.In every case, however, two ligand anions were required to form the activated state, e.g. FelI1 (TGA)g (first-order in complex), or either of the dimers I H I H2 (second-order in complex). All these species involve deprotonated R-S- groups and sulphur complexation. The data in the heterogeneous system do not agree with this: the maximum in rate is found at pH 3.8, which is reasonable only if the RCO,H/RCO,- protolytic equilibrium is considered (Kal = 3.72 x 10-4),44 but not if -SH deprotonation is required (Ka2 = 7.9 x 10-9).44 Furthermore, the basicity of -S- groups should be larger under heterogeneous and electrostatic considerations prevent the acceptance of //O - F-0-c highly charged - Fe species as dominant.Dimer sites, such as I are ' O T O \ S-C-H - Fe-S-CH, I H not excluded by this reasoning, but at high TGA concentration these species are expected to dissociate to two monomeric complexed sites. We therefore postulate that the ensuing redox reaction is of the type: 2 FeI11-O-C(0)CH2SH- 2 2 The close proximity of FelI1 centres provides FeII + R--S-S-R (4) 1 fast FeIIaq the adequate vicinity of the two organic molecules,- fulfilling the same role as the previously mentioned homogeneous precursors. It is probable that the involvement of adjacent sites may be a general reason for rate enhancement in redox reactions when more than one electron must be exchanged. A similar case was suggested to be the oxidation of hydrazine by solid barium The influence of total thioglycolic acid concentration is shown in fig.4. This influence is well modelled by the above mentioned assumption of Langmuir adsorption as a pre-equili brium.2350 Dissolution of Cobalt Ferrites 7.5 5.0 I vl N 0, . 2.5 0 0.2 0 . L 0.6 CT,-,/mol dm-3 Fig. 4. Specific rate constant k for the dissolution of Co,~,5Fe,~,,0, as a function of total thioglycolic acid concentration at 70 "C; pH,, 2.9. 0.3 m , I n 7 0.2 - v I * 0 . 1 0 1 I I I I I 2 r, 6 f/min Fig. 5. Contracting sphere plots for the dissolution by TGA 6.5 x lop2 mol dmp3 of Co,,,,Fe,,,,O, at 70 "C in presence of ferrous ion; 0, [Fe2+] = 0; a, [Fe2+] = 9.0 x mol dm-3; a, [Fe2+] = 2.5 x mol dm-3; pH,, 2.9. The role of thioglycolic acid can be performed by other complexing anions that can be engaged in an internal electron-transfer process with FelI1.These are potentially all ligands featuring ligand-to-metal charge-transfer bands, the simplest example being SCN-. Thiocyanate dissolves magnetite in acidic solutions, both in a thermal and a photochemical processes.47 Other reported examples are I-,47 oxalate17 and itr rate.^^ At high x, addition of ferrous salts supresses the lnduction period, and gives rise to a faster dissolution, see fig. 5. This is an important result that shows (a) that the reaction is not diffusion controlled, (b) the subtleties of the mechanism to be discussed below.M . A . Blesa, A . J . G. Maroto and P. J . Morando 2351 The Influence of Solid Composition on the Dissolution Rate When discussing the influence of solution variables, the possible importance of CoI1 substitution for FeII was ignored.For x d 0.6 this is consistent with the experimental results, and can be understood easily on the basis of a fast release of MI1 ions as compared to FeIII. This has been abundantly documented in the literature [see ref. (lo), (49) and (50)] and points to an essentially identical FeTrl surface in all cobalt-ferrites with x 6 0.6. Because of this reasoning, and because of the sharp change in behaviour observed at x x 0.6, it is not reasonable to attribute the onset of an induction period to different phase-transfer rates of CoII and FeII.? On the other hand, it is well known that the electronic structure of the solid suffers a similar drastic change in the same composition range, owing to the isolation of FelI ions from other FeT1 centres.26327 It is not at all obvious why the two changes should be related, but we shall offer a tentative rationale.The steps in the overall dissolution process after electron transfer to Fe"I are:51 Ferlkink --+ FelIads (Stern plane) --+ FelIbulk. ( 5 ) As long as there is a possible fast electron hopping, the intermediate FelIads can be an excellent reductant for further FeIII ions,33 and successive electron-transfer processes can give rise to an appreciable 'chain length' for the dissolution process originated from one single FerI1 reduction. This is in agreement with the recent report by Tronc et ~ l . , ~ ~ who have shown that FeII adsorbed onto magnetic iron oxides can in fact pump electrons into the colloid core within the octahedral sub-lattice of the spinel structure.Exogenous FeII can play the same role, i.e. dissolve magnetite in the presence of complexing ions through an outer-sphere heterogeneous electron transfer : Kads FeIII-TGA + FeII-TGA,, -+ - Felll-TGA..*F e I1 -TGAads FE11-TGAa9 + Fe1I1-TGAaq t - FeII-TGA. ..Felrl-TGA,,,. This is similar to the reductive dissolution of magnetite by FeII in ~ x a l i c , ~ ethylendiaminetetra-acetic6T 52 and nitrilotriaceticll acid solution^.^^ 11, 47 This interpreta- tion suggests that the overall reaction rate measured at high conversion contains, in every case, a substantial contribution from the FeIT pathway. Even the initial rate in the induction-period-free systems does not give a direct measure of the rate of reaction (4) because of the 'chain-length' effect mentioned above.Under the conditions of the present work, this effect did not give rise to noticeable localized attack, and dissolution proceeded in an essentially isotropic fashion, with only rounding off of edges. However, the possibility of a pitting attack must be taken into account; such pitting was observed by Segal and Sellers13 in the case of the dissolution of nickel ferrite by tris(picolinato)vanadium(Ir). L E T (6) M. A. B and P. J. M. are members of CONICET. Partial support through grants from SECYT-CONICET and CICPBA is gratefully acknowledged. References 1 D. L. Leussing and I. M. Kolthoff, J . Am. Chem. Soc., 1953, 75, 3904. 2 R. A. Hummel and E. B. Sandele, Anal. Chim. Acta, 1952, 7 , 308. 3 D. Bradbury, in Water Chemistry in Nuclear Reactor Systems 1 (British Nuclear Energy Society, London, 1978), p.373. i. 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