4934
Langmuir 1996, 12, 4934-4939
Reductive Dissolution of Hematite in Acidic Iodide Solutions Salvador P. Alı´, Miguel A. Blesa, Pedro J. Morando, and Alberto E. Regazzoni* Unidad de Actividad Quı´mica, Comisio´ n Nacional de Energı´a Ato´ mica, Av. del Libertador 8250, 1429-Buenos Aires, Argentina Received March 6, 1996. In Final Form: May 28, 1996X Iodide reductively dissolves hematite in acidic solutions at 353 K. Initial dissolution rates were measured under conditions of negligible back-reaction. The rate is first order in [H+] and second order with respect to iodide concentration. These results are consistent with a dissolution mechanism that involves fast equilibrated formation of inner-sphere tFe(III)sI surface complexes, fully equilibrated internal electron transfer, slow scavenging of surface coordinated I• by adsorbed I- (within the diffuse layer), and subsequent phase transfer of Fe(II) ions. The similarities between surface and solution redox chemistries are stressed, and their differences are discussed. The operation of a nonreductive parallel pathway during dissolution of iron(III) oxides by nonmetallic complexing two-electron reductants is discussed in terms of the established mechanism. The implications are tested by exploring the dissolution kinetics in iodide-thiocyanatecontaining solutions. The synergism predicted by the postulated mechanism is indeed observed and can be described in terms of an additional path in which iodide scavenges chemisorbed SCN• generated by internal electron transfer within tFe(III)sSCN surface complexes.
Introduction The dissolution of iron(III) oxides in aqueous media is a surface controlled reaction which is strongly influenced by the nature and past history of the solid phase and by the chemical composition of the swamping solution. The vast wealth of experimental information, cast in terms of the successful surface complexation description of chemical dissolution of metal oxides, has been comprehensively reviewed by Blesa et al.1 In this description of the dissolution mechanism, the strong dependence of rates upon solution composition stems from its influence on surface speciation and undergoing surface reactions. Indeed, it is very well documented that fast, fully equilibrated, complexation of surface metal centers by aggressive solutes precedes the transfer of metal ions to the aqueous phase.2-6 These surface complexes are therefore viewed as the precursors of a sequence of steps that leads to a critical complex which dissolves irreversibly through an activated state.7 Thus, specific rates of dissolution are expressed in terms of the probability of finding such a critical surface complex8 or, alternatively, in terms of the surface concentration and the reactivity pattern of the various possible surface complexes.9,10 In this work we present a detailed kinetic study of the reductive dissolution of hematite particles in acidic sodium iodide solutions, in an attempt to describe the surface X
Abstract published in Advance ACS Abstracts, August 1, 1996.
(1) Blesa, M. A.; Morando, P. J.; Regazzoni, A. E. Chemical Dissolution of Metal Oxides; CRC Press: Boca Raton, FL, 1994. (2) Segal, M. G.; Sellers, R. M. In Advances in Inorganic and Bioinorganic Mechanisms; Sykes, A. G., Ed.; Academic Press: London, 1984; Vol. 3, p 97. (3) Stone, A. T. In Geochemical Processes at Mineral Surfaces; Davis, J. A., Hayes, K. J., Eds.; ACS Symposium Series 323; American Chemical Society: Washington, DC, 1986; Chapter 21. (4) Stumm, W.; Furrer, G. In Aquatic Surface Chemistry; Stumm, W., Ed.; J. Wiley & Sons: New York, 1987; Chapter 8. (5) Stone, A. T.; Morgan, J. J. In Aquatic Surface Chemistry; Stumm, W., Ed.; J. Wiley & Sons: New York, 1987; Chapter 9. (6) Blesa, M. A.; Regazzoni, A. E.; Maroto, A. J. G. Mater. Sci. Forum 1988, 29, 31. (7) Aagaar, P.; Helgeson, H. C. Am. J. Sci. 1982, 282, 237. (8) Wieland, E.; Wehrli, B.; Stumm, W. Geochim. Cosmochim. Acta 1988, 52, 1696. (9) Hiemstra, T.; van Riemsdijk, W. H. J. Colloid Interface Sci. 1990, 136, 132. (10) Regazzoni, A. E.; Blesa, M. A. Langmuir 1991, 7, 473.
S0743-7463(96)00203-X CCC: $12.00
redox chemistry that prompts the irreversible phase transfer of iron ions in the +2 oxidation state. The similarities between solution and surface chemistries, implicit in the surface complexation approach, will be stressed. The differences will also be pointed out. The effect of thiocyanate on the dissolution kinetics is also explored. Experimental Section Hematite was synthesized as described previously;10,11 the BET specific surface area of the powder was 67.5 m2 g-1, and the average particle size was 60 nm. All solutions were prepared using analytical grade chemicals and 0.06 µS cm-1 water obtained from an E-pure Barnstead apparatus. Experiments were carried out in a 0.2 dm3 Teflon-capped Pyrex reaction vessel encasted in a jacket through which thermostated water was circulated. Dissolution experiments were performed as follows: 0.15 dm3 of a freshly prepared, O2-free, solution of desired NaI and HCl concentrations was poured into the reaction vessel and allowed to reach the working temperature (353.0 ( 0.1 K). The potential of a combined pH electrode immersed in the solution was measured and the reading used to set a Metrohm E-stat (614 Impulsomat-655 Dosimat) which was employed to add the amount of 0.25 mol dm-3 HCl that was necessary to keep a constant proton concentration throughout the kinetic run. The dissolution reaction was started by adding 0.075 g of R-Fe2O3. Addition of the solid caused a sudden drop in temperature (ca. 6 K); the original temperature was restored in less than 3 min. Aliquots of the reacting suspensions were withdrawn at different time intervals and rapidly filtered through 0.2 µm pore size cellulose nitrate membranes. The supernatants, which were particle-free, as indicated by the absence of turbidity, were stored for the assay of dissolved iron. Reacting suspensions were agitated with a 20-mm Teflon-coated stirring rod that was magnetically driven at 1000 rpm. All experiments were performed under normal laboratory illumination and carried out under a stream of N2 which was previously scrubbed through an alkaline pyrogallol solution and bubbled through a 353 K thermostated NaI solution; the concentration of NaI in the latter matched that of the reacting systems. Additional experiments were carried out in the presence of added NaSCN (1.7 mol dm-3); in these cases, [H+] was kept constant at 0.01 mol dm-3. (11) Regazzoni, A. E.; Blesa, M. A.; Maroto, A. J. G. J. Colloid Interface Sci. 1988, 122, 315.
© 1996 American Chemical Society
Dissolution of Hematite in Acidic Iodide Solutions
Langmuir, Vol. 12, No. 20, 1996 4935
Figure 1. Concentration of dissolved iron as a function of reaction time and solution composition: (b) [I-] ) 1.56 mol dm-3, [H+] ) 0.01 mol dm-3; (9) [I-] ) 1.20 mol dm-3, [H+] ) 0.01 mol dm-3; (2) [I-] ) 1.20 mol dm-3, [H+] ) 5.6 × 10-3 mol dm-3; ([) [I-] ) 0.50 mol dm-3, [H+] ) 0.01 mol dm-3; in all cases T ) 353 K.
Figure 2. Influence of iodide concentration on the specific rate of dissolution: [H+] ) 0.01 mol dm-3; T ) 353 K. The inset shows the corresponding log RS vs log [I-] plot.
The dissolved iron concentration was determined by measuring the absorbance of the iron(II) thioglycolate complex (530 ) 3770 ( 20 mol-1 dm3 cm-1) in a Shimadzu UV-210A spectrophotometer.
Results and Discussion Dissolution of hematite in acidic sodium iodide solutions at 353 K readily takes place. Figure 1 shows C(t), the concentration of dissolved iron, plotted as a function of the reaction time t, for several typical kinetic runs. In all studied cases, constant dissolution rates were rapidly attained and remained unchanged within the time span of the experiments. These results are consistent with a dissolution process that progresses, far from solubility equilibrium, with only a minor reduction of the interfacial surface area, i.e.,
dC/dt ) RSA(t)[1 - exp(∆G(t)/RT)] ≈ RSA(t0) (1) where RS is the specific rate of dissolution (in mol m-2 s-1), A(t) is the instantaneous surface area per unit volume, ∆G(t) is the time dependent driving force, R is the universal gas constant, and T is the absolute temperature. In our experiments, all collected C(t) values satisfy dC/dt ≈ RSA(t0).12,13 In doing so, we wish to place the emphasis on RS, which contains the most relevant chemical information;1 e.g., for surface-controlled dissolution reactions, the specific rate embodies the contributions arising from sites of different reactivities, viz.
RS )
∑kiΓi
(2)
where Γi is the surface density of the i sites and ki is the first-order phase transfer rate constant. All measured RS values are well below the diffusion control limit. As shown in Figure 1, RS is very sensitive toward the composition of the aqueous solution. The dependence of the specific rate of dissolution on iodide concentration at a constant [H+] (0.01 mol dm-3) is presented in Figure 2, which shows that RS increases with the concentration of iodide, the slope increasing with [I-]. The observed dependence fits very well a pseudo-second-order kinetic (12) For all measured C(t) values, calculations using ∆G°353 ) - 38.6 kJ mol-1 (for the reductive dissolution of hematite) demonstrated that the condition [1 - exp(∆G(t)/RT)] g 0.99 was fulfilled; the value for ∆G°353 was derived from the thermodynamic data compiled by Heusler and Lorenz (Heusler, K. E.; Lorenz, W. J. In Standard Potentials in Aqueous Solution; Bard, A. J., Parsons, R., Jordan, J., Eds.; Marcel Dekker: New York, 1985; p 391). (13) Even though A(t) may be estimated on the basis of isotropic contracting volume (i.e., constant penetration rate), A(t) ≈ A(t0) is the safest assumption for the very low dissolved fractions that were measured (e.g., less than 0.08).
Figure 3. Influence of proton concentration on the specific rate of dissolution; [I-] ) 1.20 mol dm-3; T ) 353 K.
law, as demonstrated by the inset in Figure 2, in which data points are replotted on a log-log scale. Kinetic orders with respect to reactant concentration larger than 1, although rare for dissolution reactions of metal oxides, have been reported earlier in the literature.14-16 The rate of dissolution of hematite in iodide solutions is also affected by solution pH. Figure 3 shows that, at a fixed [I-] (1.20 mol dm-3), RS increases linearly with increasing proton concentration. The dissolution reaction
R-Fe2O3 + 3I- + 6H+ f 2Fe2+ + I3- + 3H2O (3) may proceed through three different, parallel, pathways: (i) a direct acid path, which involves proton attack on surface Fe(III)-O oxo bonds, phase transfer of Fe(III) ions, and reduction in solution; (ii) a nonreductive iodideassisted acid path, which involves simultaneous nucleophilic-electrophilic attack by iodide and protons on surface Fe(III)-O oxo bonds, phase transfer of Fe(III)-iodide complexes, and reduction in solution; (iii) a reductive path, which involves reduction of surface Fe(III) ions, subsequent proton attack on the more labile surface Fe(II)-O oxo bonds, and phase transfer of Fe(II) ions. Since the direct acid path is not affected by iodide and the last two contributions can, in principle, be included in a single rate term (RI-), the overall dissolution rate takes the form (cf. eq 2)
RS ) RH+ + RI-
(4)
Figure 2 shows that the contribution from an iodide independent pathway is negligible for [I-] g 0.5 mol dm-3 (i.e., RI- predominates). The contribution due to the nonreductive direct acid path may, however, be important (14) Pulfer, K.; Schindler, P. W.; Westall, J. C.; Grauer, R. J. Colloid Interface Sci. 1984, 101, 554. (15) Azuma, K.; Kametani, H. Trans. AIME 1964, 230, 853. (16) Sidhu, P. S.; Gilkes, R. J.; Cornell, R. M.; Posner, A. M.; Quirk, J. P. Clays Clay Miner. 1981, 29, 269.
4936 Langmuir, Vol. 12, No. 20, 1996
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at low [I-]. In fact, the influence of iodide concentration on the rate of dissolution of hematite at [H+] ) 0.01 mol dm-3 is best fitted by
109RS/mol m-2 s-1 ) (0.01 ( 0.20) + (3.00 ( 0.18)[I-]2 (5) An attempt to measure the rate of acid dissolution of hematite in 0.01 mol dm-3 HCl only allows us to set an approximate upper limit for the nonzero ordinate, ca. 1×10-11 mol m-2 s-1, which agrees, within an order of magnitude, with the rate value that can be extrapolated from previously reported data.15-17 If this is a valid estimate, the contribution arising from the direct acid path to the RS values measured at 1.20 mol dm-3 NaI and varying acidities (Figure 3) can be safely neglected, since RH+ decreases with decreasing [H+]; note that RH+, which embodies any possible effect due to chloride ions, was reported to be a quadratic function of HCl concentration.15,16 Therefore, our data reflect essentially the influence of solution composition on the specific rate of hematite dissolution by iodide attack, i.e., RS ≈ RI-, which complies with the empirical rate law
RI- ) k[H+][I-]2
electron transfer to surface Fe(III) seems to be restricted to very strong one-electron reductants, such as the aqueous electron23 and viologen radicals;24,25 the operation of an outer-sphere mechanism during the reductive dissolution of iron(III) oxides by strongly reducing metal complexes26-29 has been recently questioned.30 Inner-sphere electron transfer in the reductive dissolution of Mn(III,IV) oxides has been thoroughly discussed by Stone.31 Therefore, to account for the mechanism of hematite dissolution by iodide attack, we postulate that the tFe(III)sOH2+‚‚‚Iouter-sphere surface complexes act solely as precursors in the fast formation of tFe(III)sI, as in an Eigen-Wilkins mechanism (see, e.g., ref 22), the inner-sphere surface complexes playing the foremost role in the dissolution process. The dissolution mechanism is accounted for by the following scheme:
tFe(III)sOH + I- + H+ a tFe(III)sI + H2O; KIint (7) III kpht
fast
tFe(III)sI 98 FeI2+ 98 Fe2+ + 1/2I2
(6)
ket
tFe(III)sI {\ } tFe(II)sI• k
Dissolution Mechanism To describe dissolution processes in terms of the surface complexation approach, it is first required to unravel the nature of the reactive surface complexes. Adsorption of iodide onto the positively charged hematite surface (at 353 K, the pH of the point of zero charge of R-Fe2O3 is ca. 7.60)18,19 must essentially lead to the formation of outersphere tFe(III)sOH2+‚‚‚I- surface complexes. However, in spite of the poor affinity of iodide for Fe(III) (iodide must be the weakest ligand among the halides and pseudohalides),20 inner-sphere tFe(III)sI surface complexes may also form. For weakly complexing reductants, such as iodide, inner-sphere surface complexation must, certainly, be determined by the interplay between outersphere electron transfer and surface ligand exchange rates. In solution, formation of aqueous Fe(H2O)5I2+ is inhibited by the faster outer-sphere reduction of Fe(H2O)63+;21 neglecting the retarding effect of Fe2+, the rate of the latter process is determined by the rate of internal electron transfer within the ion triplet I-‚‚‚Fe(H2O)63+‚‚‚I-.21 Such a trend should, however, revert in the heterogeneous R-Fe2O3/iodide solution system. Indeed, ligand exchange rates at the hematite surface are expected to be much faster than those in the analogous homogeneous systems because partial hydrolysis of Fe(H2O)63+ increases substantially the rate of water substitution.22 Moreover, if we assume that the heterogeneous and homogeneous outer-sphere reduction of Fe(III) by iodide proceeds through similar activated complexes, the structured nature of the interface would hinder the formation of the ion triplet, hence, decreasing the rate of the heterogeneous outer-sphere electron transfer. Besides, outer-sphere (17) Kametani, H.; Azuma, K. Trans. AIME 1968, 242, 1025. (18) Fokkink, L. G. J.; de Keizer, A.; Lyklema, J. J. Colloid Interface Sci. 1989, 127, 116. (19) Blesa, M. A.; Maroto, A. J. G.; Regazzoni, A. E. J. Colloid Interface Sci. 1990, 140, 287. (20) Smith, R. M.; Martell, A. E. Critical Stability Constants; Plenum Press: New York, 1976; Vol. 4. (21) Laurence, G. S.; Ellis, K. J. J. Chem. Soc., Dalton Trans. 1972, 2229. (22) Margerum, D. W.; Cayley, G. R.; Weatherburn, D. C.; Pagenkopf, G. K. In Coordination Chemistry; Martell, A. E., Ed.; ACS Monograph 174; American Chemical Society: Washington, DC, 1978; Vol. 2, Chapter 1.
(8) (9)
-et
kscav
tFe(II)sI• + I-(ads) + H2O 98 tFe(II)sOH2 + I2•- (10) II kpht
tFe(II)sOH2 98 Fe2+
(11)
The diiodine radical (E°(I2/I2•-) ) 0.11 V)32 may further react with an adjacent surface iron(III) ion, irrespective of its protonation state, fast
tFe(III)sOH + I2•-(ads) + H+ 98 tFe(II)sOH2 + I2 (12) or, more likely, decompose very rapidly in solution,33 e.g.
I2•- + I2•- f I3- + I-; k ) 3.9 × 109 mol-1 dm3 s-1 (13) I2
•-
2+
+ Fe
2+
f FeI
+I ; -
k ) 3.6 × 106 mol-1 dm3 s-1 (13a) (23) Buxton, G. V.; Rhodes, T.; Sellers, R. M. Nature (London) 1982, 295, 538. (24) Mulvaney, P.; Swayambunathan, V.; Grieser, F.; Meisel, D. J. Chem. Phys. 1988, 92, 6732. (25) Mulvaney, P.; Cooper, R.; Grieser, F.; Meisel, D. Langmuir 1988, 4, 1206. (26) Segal, M. G.; Sellers, R. M. J. Chem. Soc., Faraday Trans. 1 1982, 82, 1149. (27) Baumgartner, E. C.; Blesa, M. A.; Marinovich, H. A.; Maroto, A. J. G. Inorg. Chem. 1983, 22, 2224. (28) Blesa, M. A.; Marinovich, H. A.; Baumgartner, E. C.; Maroto, A. J. G. Inorg. Chem. 1987, 26, 3713. (29) Borghi, E. B.; Regazzoni, A. E.; Maroto, A. J. G.; Blesa, M. A. J. Colloid Interface Sci. 1989, 130, 299. (30) Hering, J. G.; Stumm, W. In Mineral-Water Interface Geochemistry; Hochella, M. F., Jr., White, A. F., Eds.; Reviews in Mineralogy; Mineralogical Society of America: Washington, DC, 1990; Vol. 23, Chapter 11. (31) Stone, A. T. Environ. Sci. Technol. 1987, 21, 979. (32) Wardman, P. J. Phys. Chem. Ref. Data 1989, 18, 1637. (33) Neta, P.; Huie, R. E.; Ross, A. B. J. Phys. Chem. Ref. Data 1988, 17, 1027.
Dissolution of Hematite in Acidic Iodide Solutions
Langmuir, Vol. 12, No. 20, 1996 4937
In the kinetic scheme above, eq 7 depicts the chemisorption pre-equilibrium, and recalling that RI- embodies two rate contributions, i.e.
RI- ) RIII + RII
(14)
where the subscripts indicate the oxidation state of dissolving iron ions, eqs 8 and 9-11 represent, respectively, the nonreductive and the reductive dissolution pathways. At the steady state, eq 8 leads to III RIII ) kpht {tFe(III)sI}
(15)
where the braces denote surface concentration. Internal electron transfer within tFe(III)sI (eq 9) must be fully equilibrated because the fast tFe(II)sI•(0) f tFe(III)sI(-I) recombinationsa process which is downhill in energy (E°(I•/I-) ) 1.3 V)32 and, certainly, kinetically favored by surface cage effects10,34srenders k-et . kscav[I-]. Thus, at the steady state, eqs 9-11 yield II {tFe(II)sOH2} ) kscavKet{tFe(III)sI}[I-] RII ) kpht
(16) and the overall dissolution rate, given by eq 14, transforms into III + - 2 RI- ) KIint - {tFe(III)sOH}[H ](kpht[I ] + kscavKet[I ] )
(17) int -3 mol-2 dm6 (see Since KIint - is very low, i.e. KI- e 10 below), the surface concentration of uncomplexed sites remains virtually constant within the spanned experimental conditions (cf. also refs 8 and 11). Therefore, eq 17 reduces to the empirical rate law 6 only if the reductive III /[I-] , kscavKet). The experipath predominates (i.e., kpht mental evidence demonstrates the operation of a purely reductive dissolution mechanism and supports the implicit assumption leading to eq 16, i.e., the iodide ions involved in the scavenging of coordinated I• (eq 10) must be adsorbed within the diffuse layer (i.e., at the δ-plane); in fact, those adsorbed in the form of inner-sphere surface complexes must migrate against an attractive electrostatic micropotential. Note that the linear dependence of the scavenging rate on [I-] stems from the near constancy (and very low values) attained by Ψδ under our experimental conditions. Thorough descriptions of the electrical double layer formed at metal oxide/aqueous solution interfaces may be found elsewhere (see e.g. refs 1, 6, and 11 and references therein). As discussed above, different mechanisms operate during the heterogeneous and homogeneous reduction of iron(III) by iodide. Notwithstanding, the stoichiometry of the activated complexes remains the same; i.e., both processes obey second-order kinetics on iodide concentration. The kinetic scheme proposed for the reduction of surface Fe(III) (eqs 7, 9, and 10), coupled with eqs 12 and/ or 13, is formally identical to that advanced by Fudge and Sykes35 and later reexamined by Laurence and Ellis21 for the homogeneous reaction. Equations 9 and 10 may be represented, in principle, as a single reaction step. However, such a representation would imply rate control by electron transfer. Instead, we decided to write them separately to emphasize that the rate of the heterogeneous process is determined by the slow scavenging of the
(34) Borghi, E. B.; Morando, P. J.; Blesa, M. A. Langmuir 1991, 7, 1652. (35) Fudge, A. J.; Sykes, K. W. J. Chem. Soc. 1952, 119.
Figure 4. Influence of iodide concentration on the specific rate of dissolution of hematite in iodide-thiocyanate-containing solutions: [SCN-] ) 1.7 mol dm-3; [H+] ) 0.01 mol dm-3; T ) 353 K. The dotted line is the sum RI- + RSCN- (see text).
chemisorbed iodine atom, i.e., by the slow decomposition of the tFe(II)sI• successor complex. The noticeable sluggishness of the surface reaction (judged by the rate of dissolution) as compared to the very fast homogeneous process must be traced back to the fast tFe(II)sI• f tFe(III)sI recombination, which plays a role analogous to the well-known retarding effect of Fe2+ in the homogeneous reaction. In principle, dissolution of iron(III) oxides by nonmetallic two-electron reductants, such as iodide, thiocyanate, and mercaptocarboxylic acids, should follow the same general mechanism. In contrast with iodide, dissolution by the latter reagents obeys a Langmuir-Hinshelwood kinetic law.10,34 This behavior cannot be solely attributed to the much larger affinities of SCN- and mercaptocarboxylates for tFe(III), since, at saturation coverage, eq 16 predicts a first-order dependence on reductant concentration. Hence, compliance with a Langmuir-Hinshelwood rate law would reveal the predominance of the nonreductive ligand-assisted acid path. Indeed, a very important contribution of the RIII path has been diagnosed for the latter systems. Although the coupled fast homogeneous reduction of dissolved Fe(III)L complexes (cf. eq 8) precluded a meaningful assessment of RIII/RII ratios, in the case of thiocyanate,10 the data seem to support III kpht /[SCN-] . kscavKet (cf. eq 17). The different dissolution behaviors of hematite in iodide and in thiocyanate solutions can be easily interpreted. III III For simplicity, we shall assume that kpht,I - ∼ kpht,SCN(actually, they may differ substantially). Since the rates of the homogeneous dimerization, X- + X• f X2•- (X- ) halides and pseudohalides), are similar,36 we can further assume that the respective kscav values are comparable. Then, Ket,SCN- , Ket,I-. This trend parallels that of the E°(X•/X-) potentials (cf. ref 32), and if a linear free energy relationship (LFER) is obeyed, the ratio Ket,SCN-/Ket,I- would be on the order of 10-6. Thus, the minor RII contribution in the case of dissolution by thiocyanate reflects the more unfavorable thermodynamic requirement for electron transfer to occur. Dissolution by iodide is, however, slower; overall dissolution rates in thiocyanate solutions are about twice or even larger, depending on concentration (see Figures 2 and 4; cf. also ref 10). This is a result of the much higher affinity of thiocyanate. Note that, on the basis of the well-known LFER between surface and solution complexation constants,37-39 extrapolation of stability constants for Fe(III)-X aqueous complexes20 (36) Baxendale, J. H.; Bevan, P. L. T.; Stott, D. A. Trans. Faraday Soc. 1968, 64, 2389. (37) Stumm, W.; Kummert, R.; Sigg, L. Croat. Chem. Acta 1980, 53, 291. (38) Sigg, L.; Stumm, W. Colloids Surf. 1981, 2, 101. (39) Dzombak, D. A.; Morel, F. M. M. Surface Complexation Modeling. Hydrous Ferric Oxide; J. Wiley & Sons: New York, 1990.
4938 Langmuir, Vol. 12, No. 20, 1996
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int int yields 104 as a rough estimate for KSCN -/KI- (the actual ratio may be even larger; cf. ref 21). This makes III int III int kpht,SCN - × KSCN- . kpht,I- × KI- (cf. eq 17). In spite of the somewhat different description of the SCN- chemisorption pre-equilibrium presented in ref 10, KIint - can be estimated to be e10-3 mol-2 dm6. Dissolution of magnetite by mercaptoacetic, 3-mercaptopropionic, and mercaptosuccinic acids is a more complex III values are greatly inprocess;34 in these systems, kpht creased due to the trans-effect that is brought about by ring-closure, and dissolution rates decrease with increasing ligand affinity. Even though comparison with the studied system is by no means straightforward, the absence of a second-order rate term could be interpreted on the basis of the above rationale, since, despite the favorable one-electron potential (E°(RS•/RS-) ca. 0.7 V),32 the exceedingly low concentration of RS- at the pH of interest40 should make scavenging of the chemisorbed radical very slow. This, however, does not necessarily imply that the RIII path predominates, because the expectedly large Ket value may render direct phase transfer of tFe(II)sS•R the prevailing dissolution path; likewise, reductive dissolution of magnetite and hematite by ascorbate follows a Langmuir-Hinshelwood dependence with reductant concentration.41,42 A further complexity of these systems is the parallel reductive path that is triggered by the accumulation of dissolved Fe(II), which, at a certain conversion degree, dominates the overall kinetics.
Synergism in Iodide-Thiocyanate-Containing Solutions The proposed mechanism (eqs 7-11) suggests that overall dissolution rates can be greatly enhanced by the addition of adequate X• scavengers, thus increasing the contribution due to RII by triggering a parallel reductive path in which tFe(II)sX- surface complexes are rapidly formed and transferred to the aqueous phase; in the limit, the rate of this reductive path may be controlled by electron transfer itself, i.e., k′scav[scav] . k-et, or even by the formation of the precursor tFe(III)sX surface complex. Noting the thermodynamic feasibility of reaction
I- + SCN• a I• + SCN-; K ) 3.69 × 104 (18) (see ref 43), the rate of hematite dissolution by thiocyanate-iodide mixtures was measured to test this hypothesis. The results presented in Figure 4 demonstrate that hematite does, indeed, dissolve faster in solutions containing both iodide and thiocyanate anions; the rates are much larger than the sum of the individual contributions, RI- + RSCN-, which is depicted by the dotted line. The empirical kinetic order with respect to [I-] changes in the presence of thiocyanate, thus providing strong evidence of a marked change in mechanism. In principle, the overall dissolution rate can be expressed as
RS ) RI- + RSCN- + Rsyn
(19)
where Rsyn represents the increase in rate produced by synergism. The individual rate contributions are not necessarily those measured in the separate experiments, (40) Martell, A. E.; Smith, R. M. Critical Stability Constants; Plenum Press: New York, 1977; Vol. 3. (41) dos Santos Afonso, M.; Morando, P. J.; Blesa, M. A.; Banwart, S.; Stumm, W. J. Colloid Interface Sci. 1989, 131, 567. (42) Banwart, S.; Davies, S.; Stumm, W. Colloid Surf. 1989, 39, 303. (43) Scho¨nesho¨fer, M.; Henglein, A. Ber. Bunsen-Ges. Phys. Chem. 1970, 74, 393.
for competitive adsorption affects both; owing to the large int int KSCN ratio, RI- should decrease substantially -/KIwhereas RSCN- should remain essentially the same. Thus, at constant [SCN-] and [H+], any contribution from Rsyn must reflect itself as an increase in RS. Since, SCN- may also scavenge I•, two parallel steps must be considered to be possibly involved in the observed synergistic effect: kscav,1
tFe(II)sSCN• + I-(ads) 98 tFe(II)sSCN- + I• (20) kscav,2
tFe(II)sI• + SCN-(ads) 98 tFe(II)sI- + SCN• (21) The radicals rapidly evolve in solution, and the surface complexes dissolve. Thus, at the steady state, Rsyn is given by
Rsyn ) kscav,1{tFe(II)sSCN•}[I-] + kscav,2{tFe(II)sI•}[SCN-] ) kscav,1Ket,SCN- ×
(
{tFe(III)sSCN}[I-] 1 +
kscav,2Ket,I-KIint int kscav,1Ket,SCN-KSCN -
)
(22)
which, at a constant thiocyanate concentration and [H+], reduces to
Rsyn ) k′[I-]
(23)
in excellent agreement with the observed kinetic order (Figure 4). Whether step 20 or 21 predominates remains unclear; a rough estimation of kscav values based on the ratio k18/k-18 would favor step 20. Nevertheless, the role of thiocyanate should be viewed as that of a catalyst, the reaction stoichiometry being that of the reductive dissolution of hematite by iodide. The alternative view, invoking k′scav,1
tFe(II)sSCN• + I-(ads) + H2O 98 tFe(II)sOH2 + ISCN•- (24) k′scav,2
tFe(II)sI• + SCN-(ads) + H2O 98 tFe(II)sOH2 + ISCN•- (25) is kinetically indistinguishable. Scavenging of coordinated radicals may not be restricted to nonmetallic reductants. In principle, the well-documented rate acceleration produced by leached Fe(II) during dissolution of iron(III) oxides by mercaptocarboxylates may also be interpreted in terms of the above rationale. It is worth stressing, however, that due account of solution and surface speciation ought to be taken prior to reaching definite conclusions; the complex multiplicity of reaction pathways leading to synergism has been nicely illustrated previously.41,44 General Conclusion The studied case is another example of the universality of the basic concepts of coordination chemistry and (44) Rueda, E. H.; Ballesteros, M. C.; Grassi, R. L.; Blesa, M. A. Clays Clay Miner. 1992, 40, 575.
Dissolution of Hematite in Acidic Iodide Solutions
provides a nice illustration of the well-known parallelism between solution and surface chemistries. It further illustrates that analogous heterogeneous and homogeneous reactions may proceed through different mechanisms and, yet, the stoichiometry of the activated complex be the same. This finding indicates that straightforward extrapolation from solution chemistry may not be necessarily guaranteed, even in the cases where the similarities
Langmuir, Vol. 12, No. 20, 1996 4939
between surface and solution chemistries are not blurred by the manifold of dissolution pathways. Acknowledgment. MAB, PJM, and AER are members of CONICET. Partial support by Fundacio´n Antorchas is gratefully acknowledged. LA960203U
