Langmuir 1991, 7, 1652-1659
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Dissolution of Magnetite by Mercaptocarboxylic Acids E.B.Borghi, P. J. Morando, and M. A. Blesa* Departamento Quimica de Reactores, Comisibn Nacional de Energia Atbmica, Avda. del Libertador 8250, 1429 Buenos Aires, Argentina Received June 7,1990. In Final Form: February 11 , 1991 Mercaptocarboxylicacids dissolve magnetite by creating reactive surface complexes in which the ligand is bound to surface iron(II1) centers via both 400-and -S- groups. These surface complexes dissolve rather rapidly by a mechanism that probably involves both sequential redox decomposition-dissolution and dissolution-redox decomposition. In either case the causative factor of rate enhancement is the weakeningof Fe-0-Fe bonds caused by the ligand. Very early in the dissolution process a further catalytic pathway is established by dissolved FeLL complexes. The dependence of the rate of all these processes on solution variables (pH, ligand nature and concentration) is essentially determined by the changes brought about in the number of surfacecomplexes createdby adsorption. The reactivity of these complexes is much less sensitive to the mentioned variables. Hydrogen peroxide addition gives rise to a complex sequence of reactions that involves free radicals and leads to the oxidation of the ligand as the main stoichiometric reaction.
Introduction The kinetics and mechanisms of reactions of species dissolved in liquids can usually be reasonably well described in terms of solvated molecules that interact with a continuum solvent. Solids, on the other hand, cannot be resolved in general into small discrete molecular units, and the attempts to describe their reactivity usually rely on the idea of "(re)active sites" associated with some kind of faulty structure. It is not surprising that a further step should be taken by invoking concepts such as "surface states" and "surface complexes" so as to more closely approach the idea of small discrete units, rather decoupled from the rest of the solid framework. Although it is clear that "surface complexes" are not molecules, many characteristics of the interaction between metal oxide particles and dissolved complexing ligands can be rationalized on the basis of treating isolated metal centersbound to the ligand as "semicomplexes".1?2 By now, an extensiveinterfacial coordination chemistry is beginning to emerge from the studies of chemisorption of ligands onto metal 0xides.3-5 In particular, the similarities and differencesbetween surfaceand bulkchemistry are an important research front. This paper explores the heterogeneous chemistry of a series of thiols with Fe(II1) in the magnetite (FesO4) surface. Because of the high rates of dissolution,the experimentaldata used to construe a picture of the involved heterogeneous complexation chemistry are essentially rates of dissolution. In aqueous solution, Fe(II1) and thiocarboxylic acids form blue FeIIIL+species (stability constant K ) , in which the metal ion is probably bound to -S-R and -C(O)O-. In acidic media, the conditional stability constants, K', of FenxL+are not high because of the two successive protonations of L2- (for H2L pKa1 2-4, pKa2 10-11). Rather modest changes in K (e.g. from 3.1 X 10'3 for thio-
-
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(1) Valverde, N. Ber. Bunsengee. Phye. Chem. 1976,80, 333. (2) Schindler, P. W.; Stu", W. Aquatic Surface Chemistry;Stumm, W., Ed.;John Wiley & Sons: New York, 1987; Chapter 4. (3) Blese, M. A.; Regazzoni, A. E.; Maroto, A. J. G. Mater. Sci. Forum 1988, 29, 31. (4) Hingston, F. J. Aduorption of Inorganics at Solid-Liquid Interfaces;Anderson, M. A., Rubin, A. J., Eds.; Ann Arbor Science Publishers: Ann Arbor, MI, 1981; Chapter 2. (5) Hohl,H.;Sigg, L.; Stumm, W. Particulates in Water;Kavanangh, M. C., Leckie, J. O., Eds.; Advances in Chemistry Series 189; American Chemical Society: Washington, DC, 1980, Chapter 1.
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glycollate to 1.4 X 10" for penicillamine dianion)s may bring important changes in the speciation: for the above cases at pH 2, K' values are respectively 10.'and 102; when [LIT, the total uncomplexed ligand concentration, is ([FeL+]/ [Fe3+])is therefore 10 in the first case and 0.1 in the second. It was considered likely, therefore, that the extent of chemisorption could also change appreciably when rather small changes are introduced in the ligand molecule. In aqueous solution, these complexes undergo a fast internal redox reaction that yields Fe(I1) and disulfide (reaction 1)."lo The rate of (1) decreases with increasing acidity, making the blue colors observable directly only at very low pH values. FeL+
-
Fen + '/,RS-SR'-
(1)
The heterogeneous chemistry of thiocarboxylates and Fe(II1)-containing oxides has been explored to a much lesser extent.11-14 The overall result of the interaction is the reductive dissolution of the oxide, a fact extensively used in the analysis of iron oxides with thioglycolic acid (TGA).I6 Our earlier studies of the systems magnetite/ TGA12 and ferrites/TGA13 demonstrated the operation of a dissolution mechanism involvingadsorption, internal redox reaction, and phase transfer of Fe(I1) to complete dissolution.
Experimental Section Materials. Magnetite was prepared by following a standard procedure.16 The sample used in these experimentswas the same (6) Baiocchi, C.; Mentaeti, E.; Arselli, P. TransitionMet. Chem. 1983, 8,40. (7) bussing, D. L.; Kolthoff, I. M. J. Am. Chem. SOC.1969,76,3904. (8) Tanaka, N.; Kolthoff, I. M.; Stricks, W. J. Am. Chem. SOC.1966, 77,2004. (9) Jameson,R. F.; Lmert, W.; Techmkowitz,A.; Gutmann, V. J. Chem. SOC.,Dalton Trans. 1988,943. (10) Jameson, R. F.; Lmert,W.; Tschinkowitz, A. J. Chem. Soc.,Dalton Trans. 1988,2109. (11) Bradbury, D. Water Chemiutry of Nuclear Reactor S y s t e m ; British Nuclear Energy Society: London, 1978; Vol. 1, p 373. (12) Baumgartner, E.; Bleea, M. A,; Maroto, A. J. G. J. Chem. SOC., Dalton Trans. 1982, 1649. (13) Blesa, M. A.; Maroto, A. J. G.;Morando, P. J. J. Chem. SOC., Faraday Trans. 1 1986,82, 2345. (14) Waite, T. D.; Torikov, A,; Smith, J. D. J. Colloid Interface Sci. 1986,112,412. (15) Hummel, R. A.; Sandell, E. B. A d . Chim. Acta 1962, 7, 308. (16) Regazzoni, A. E.; Urrutia, G. A.; Bleea, M. A.; Maroto, A. J. G. J. Inorg. Nucl. Chem. 1981,43,1489.
0 1991 American Chemical Society
Langmuir, Vol. 7, No. 8, 1991 1653
Dissolution of Magnetite by Mercaptocarboxylic Acids as used previous1y;"Ja the average length of the cuboctahedral particles was 0.27 pm and the specific BET surface area of the solid was 9.7 mz gl. The mercaptocarboxylicacidsused wereanalyticalgradeTGA, mercaptoeuccinicacid (MSA),3-mercaptopropionicacid (MPA), and cysteine (Cys). Kinetics of Dissolution. Kinetic experiments were performed in a cylindrical reaction vessel provided with a water jacket and stirred magnetically; 0.1 dms of mercaptocarboxylic acids of adequate concentration and pH were thermostated at 55.0 & 0.1 "C, and dissolution reactions were started through the addition of 20 mg of magnetite. In some experiments the initial solutionsalsocontainedH202or Fe(II)(asFe(NHMSO&6H20). Periodical samples were taken, quenched by dilution, and filtered immediately through a Sartorious membrane (0.45 pm pore size). Total Fe was determined photometrically, using alkaline thioglycollate solutions. All experiments were carried out under a Nz atmosphere from which traces of 0 2 had been removed by scrubbing the gas through an alkaline pyrogallol solution. The pH was monitored during the course of reaction with a combined glass electrode immersed in the suspension. In no case were pH drifte larger than 0.005 unit observed. The response of the electrode was periodically checked to ensure that the suspended particles did not affect it. All points presented in the results section are the average of at least duplicate measurementsthat differed by no more than 15 96. Ionicstrength was variable and was governed essentiallyby the ionized fraction of the ligand. Spectral Characterization of Iron(II1) in Reactive Mixtures. A reacting suspension with large amounts of solid in 0.6 M MSA was prepared in a 1-cm optical cell. Deoxygenated solutions were used to fill the total volume of the cell, that was tightly stoppered except in experimentsdesigned to explore the influence of dissolved oxygen. Sequential visible absorption spectra were then recorded in a Hewlett-Packard 8452A diode array spectrophotometer at room temperature. Light scattering effects, although clearly observed in the shifts of the baselines, did not interferein the detection of a broad band centered at 550 nm. Ligand and temperature were chosen to decrease the rate and to permit measurements during longer (-30 min) times. Electron Microscopy. Occasional observation of partially reacted particles was performed in a Philips SEM apparatus.
Results Stoichiometry of the Reaction. Except for the case of cysteine, total dissolution can be achieved, and the final stoichiometry of the reaction can be written as
+
-
6H' Fe,O, + 2H2L 3Fe2' + (LH), + 4H20 (2) Cysteine is much less reactive toward magnetite; at 55 O C , pH 3, a 0.1 M cysteine solution dissolves only 20% of the solid after 2 h of contact time. In the final solution no Fe(1II) is present because of the fast reduction by H2L. The oxidation product (LH)2was not determined experimentally. During the course of reaction, the pH does not change appreciably because of the buffering by H2L/HL-. In the presence of hydrogen peroxide, the reaction
-
H20, + 2H2L 2H20+ (LH), (3) prevails until H202 or H2L is consumed. In the former case, reaction 2 follows; in the latter case, no further reaction is observed and dissolution is totally arrested. Although Fe(II1) is not present in the final solution, it is formed in intermediate stages of reaction and later consumed. Figure 1shows the sequential visible absorption spectraof a suspension containing magnetite and MSA at pH 4 a t room temperature. The broad maximum (17) Regazzoni,A. E.;Bleea, M.A.; Maroto,A. J. G.J. Colloidlnterjace Sci. lSl,91,560. (18) Reaazmni. A. E. Ph.D. Thesis, Universidad Naciond de Tucu-
m&, Argeitina, i9u.
$0.3 "'
+---< ? - - - - - - -
J L
4 00
600
5W Wavelength l ~ m l
700
800
Figure 1. Sequential visible spectra of aqueous suspensionsof FeaOl in MSA (0.6 mol dm+) solutions: a, 0; b, 4; c, 8; d, 14; e, 18; f, 22; g, 26; h, 28; i, 30 min. Spectrum h is an unstured suspension. Spectrum i is in the presence of dissolved oxygen.
Time (minl
Figure 2. Plots of magnetite dissolved fraction as a function of time for various experimental conditions: ( 0 )[MPA] = 0.2 mol dm-3, pH = 4.08; ( 0 ) [MSA] = 0.1 mol dm+, pH = 3.80; (A) [MSA] = 0.1 mol dm-3, pH = 3.85, [Fe2+]= 1.3 X mol dm+; ( 0 )[TGA] = 0.033 mol dm+, pH = 4.00. centered at 550nm corresponds to FenLMSA2-complexes; the shift in the maxima from reported values (640 nm)6 and the broadening of the band can be attributed to the dispersion caused by the particles; adsorbed complexes may also contribute to the observed effects. Although the ratio [Fe(I11)]/[Fe]~was not determined in these experiments, the comparison with the kinetic runs demonstrates that only a minor fraction of dissolved iron is present as Fe(II1). This minor fraction however implies a fast production that compensates the fast redox decomposition of Fe(II1) in thiocarboxylate solutions at pH 4.9 In other words, the overall stoichiometry includes a t least a very important contribution from a path formed by the two consecutive reactions (4 and 5). Fe,O,
+ 2H2L + 4H+
-
2FeL+(aq)+ 2H'
2FeL'
-
+ Fe2' + 4H,O
2Fe2++ (LH),
(4) (5)
As shall be discussed later, reaction 4 is highly sensitive to catalysis by Fe2+. Kinetics of Dissolution. The plots of dissolved fraction, f , vs time are in general sigmoidal, as in Figure 2. The morphology of the acceleratory stage depends on the nature of the ligand. It is also very much affected by the presence of HzOzand, to a lesser extent, by the addition of ferrous salts. TGA exhibits a relatively short acceleratory stage that justifies our previous usage of a contracting geometry rate law to describe the kinetics of d i s s ~ l u t i o n . ~ ~ J ~
Borghi et al.
1654 Langmuir, Vol. 7, No.8, 1991
Such a description is no longer valid for the whole series of ligands. The length of the acceleratory stage increases in the sequence TGA < MPA < MSA (
