1206
Langmuir 1988,4, 1206-1211
pears to be limited by slow kinetics since a substantial amount of surface hydroxyl is observed experimentally. The experimentally observed dissociation products of H20 on Ti are thus supported by the thermodynamics of the proposed dissociation mechanism. Hydrogen on the surface of Ti has a tendency to diffuse either into or out of the sample depending on the surface oxygen concentration and sample temperat~re.~'The small initial increase (Figure 2) in surface hydrogen for low H 2 0 exposures (2.5 is substantially lower than the yield expected from the complete transfer of electrons to the particles. Adsorption isotherms for Fe2+on these colloids indicate that adsorption of the reduced iron can only account for a small fraction of the reduction yields obtained. This observation, as well as measurements of the amounts of Fe2+that can be removed from the surface after the reduction, leads to the conclusion that a substantial fraction of the electrons transferred to the colloid is able to migrate into the particle's interior. Analysis of the thermodynamics of the system indicates that magnetite is an expected product of the partial reduction. The factors that control surface trapping and adsorption vs migration into the bulk are discussed.
Introduction Iron oxides are often the primary product from the oxidation of iron and steel surfaces. Corrosion processes on these surfaces result in the formation of oxide films and/or colloidal iron oxides. Further interest in these oxides arises from their frequent use as catalysts,pigments, carriers, etc. Consequently, the mechanisms which lead to the formation of these iron oxide colloids have received a great deal of attenti0n.l Exact conditions necessary to produce suspensions of iron oxide particles of well-defined sizes, morphology, and crystal structure have been developed and reviewed by Matijevic and co-workers.2 In addition, a considerable body of information exists on the dissolution of various oxides, in particular magnetite3 and ferrites. These studies have examined the role and efficacy of various complexing and reducing agents in oxide dissoluUniversity of Melbourne.
* Argonne National Laboratory.
t i ~ n . ~More , ~ recently, the need for a fast and economical method for the removal of metal oxide deposits from the primary coolant circuits of water-cooled nuclear reactors and the use of some metal oxides in energy conversion systems have prompted renewed interest in the general concepts of reductive dissolution.6 These studies involved (1)Dousma, J.; De Bruyn, P. L. J. Colloid Interface Sci. 1978,64,154.
(2) (a) Matijevic, E.; Scheiner, P. J. Colloid Interface Sci. 1978, 63, 509. (b) Matijevic, E.Acc. Chem. Res. 1981, 14, 22. (3)(a) Blesa, M. A.; Marinovich, H. A.; Baumgartner, E. C.; Maroto, A. J. G. Inorg. Chem. 1987,26, 3713. (b) Blesa, M. A,; Borghi, E. B.; Maroto, A. J. G.; Reggazoni, A. E. J. Colloid Interface Sci. 1984,98,295. (4) (a) Allen, G. C.; Kirby, C.; Sellers, R. M.; J. Chem. SOC.,Faraday Tram. 1,1988,84,355. (b) Gorichev, I. G.; Kipriyanov, N. A. Russ. Chem. Reo. 1984,53, 1039. (c) Stone, A. T.; Morgan, J. J. In Aquatic Surface Chemistry-Chemical Processes a t the Particle- Water Interface; Stumm, W., Ed.; Wiley: New York, 1987. (d) Blesa, M. A.; Maroto, A. J. G. In Proceedings, International Joint Meeting ANS-CNA on Decontamination of Nuclear Facilities (Keynote Addresses), Niagara Falls, Canada, 1982. (5) Frenier, W. W.; Growcock, F. B. Corrosion 1984, 40, 663.
0143-1~63/88/2404-~206$01.50/00 1988 American Chemical Society
Langmuir, Vol. 4, No. 5, 1988 1207
Reductive Dissolution of Iron(III) Oxides
comparison between one- and two-electron reductants, effect of chelating agents in enhancing the dissolution rate, and comparison between various iron oxides and spinels.'* Buxton et al.1° have recently studied the reductive dissolution of iron oxides by radiolytically produced reducing free radicals. Among their observations, they noticed that reduced Fe(II1) dissolved preferentially over original lattice Fe(I1) ions in Fe304. In an effort to enhance our understanding of the mechanism of reductive dissolution of iron(II1) oxides, we have studied the fate of electrons transferred to colloidal particles of these oxides using radiolytically produced viologen radicals as reductants. The advantage of using viologens as electron-transfer agents is that the mechanism of reduction is no longer masked by adsorption and the formation of transitory surface complexes, which may obscure mechanistic details of the dissolution process. In the present report, we focus on the products of the reduction process in steady-state radiolysis and transpose these observations with thermodynamic predictions of the oxides' stability.
Experimental Section Preparation and Characterization. cu-FezO3 (hematite) is a calcined ferric oxide (purchased from B.D.H.) which gave very clean X-ray diffraction powder patterns characteristic of the pure hematite phase. The BET surface area of the dry powder was 6.9 m2/g. a-FeOOH (goethite) was prepared by the method of Atkinson et al.;" 800 mL of vigorously stirred solution of 0.156 M Fe(N03)3was heated to 60 "C, and 200 mL of a solution of 2.5 M potassium hydroxide was added quickly. The optimum ratio of Fe/OH- of 0.25 has to be maintained to obtain this iron oxyhydroxide. The brown precipitate which forms originally turns into the ochre color characteristic of goethite upon aging for 24 h a t 60 "C. This precipitate was washed with water to remove excess nitrate and dried in air. Surface area, determined by BET, was 18.6 m2/g. The X-ray diffraction powder pattern revealed it to be pure goethite with no detectable amount of magnetite or hematite. To obtain colloids of smaller sizes, Sorum's method was adopted.lZ HzO (450 mL) was heated until vigorously boiling; 50 mL of freshly prepared 0.02 M FeC1, solution was added at a rate of two drops per second. The sol rapidly turned golden brown and finally deep red. After boiling for a further 5 min, the sol was allowed to cool to room temperature and then dialyzed for 48 h against HC104 at pH 3.5. This preparation yielded particles of 50 i 10 A diameter, as determined by TEM. The point of zero charge (pzc) for the particles of this sol was found to be 8.3 by electrophoresis. Electron diffraction patterns indicated a-Fe203. Analysis by the o-phenanthroline method13 M, [Fe2+(aq)]= 3 X lo4 M, and yielded [ F e l ~= 2.05 X [Fe3+(aq)]= 8 X lo-' M for this sol. The aggregation number for the sol, calculated from the radius of the particles and the density of hematite, p = 5.24 g/cm3, was n = 4rNopr3/3Mw= 1.3 X lo3 molecules/particle. Similarly, the calculated surface area of the sol was A = 3/(pr) = 230 m2/g. The transparent sol was used without any stabilizer added. No changes in the absorption spectra of the sol could be observed 2 weeks after preparation. This sol coagulated upon extended dialysis against HzO but was (6) Segal, M. G.;Sellers, R. M. In Adu. Inorg. Bioinorg. Mech. 1984, 3, 97. (7) Segal, M. G.;Sellers, R. M. J . Chem. Soc., Chem. Commun. 1980, 991. (8)Bradbury, D. Water Chemistry of Nuclear Reactor Systems; British Nuclear Energy Society: London, 1978; p 373. (9) Segal, M. G.;Sellers, R. M. J. Chem. Soc., Faraday Trans. 1 1982, 78, 1149. (10) Buxton, G.V.; Rhodes, T.; Sellers, R. M. Nature (London) 1982, 295,583; J . Chem. Soc., Faraday Trans. I 1983, 79, 2961. (11) Atkinson, R. J.; Posner, A. M.; Quirk, J . P. J.Phys. Chem. 1967, 71, 550. (12) Sorum, C. H. J . Am. Chem. SOC.1928,50, 1263. (13) Vogel, A. I. Quantitatiue Inorganic Analysis; Longmans: London, 1953.
%
5
0
a .
"L
1
0' 0.000
0.001
0.002
F e p , 11M
Figure 1. G(Fe2+)as a function of [Fez03]for the transparent sol. Total dose 10.8 h a d , pH 3.8, measured 0.5 h after irradiation. All solutions were N20 saturated and contained 0.1 M 2-propanol M V. and 2 X stable when dialyzed against acid. Some of the experiments were also repeated with transparent sols prepared according to the method of Dimitrijevic et al.14 and gave similar results to those obtained with Sorum's sols. Irradiation Procedures. Suspensions of the iron oxides were prepared with a predetermined amount of the oxide solution containing 0.1 M 2-propanol and 2 X M N,"-bis((su1fonatooxy)propyl)-4,4'-bipyridiniurn (denoted V; synthesized according to literature procedures15)at the desired pH. The solution was sonicated for 10 min and then bubbled with NzO by using the syringe technique.16 (NzOwas used to deoxygenate the solutions and to scavenge the hydrated electrons produced by radiolysis.) Slight settling of the suspensions did occur during radiolysis, but this did not affect G(Fe2+). The same procedure was used to saturate the transparent sols, except sonication was not required. Steady-state irradiation was carried out in a 6oCoy-ray source at 18had/h. After the irradiation, samples were filtered through millipore filters (pore size 0.01 pm) and analyzed for Fez+by using the o-phenanthroline method. Ferric ions were determined following reduction with hydroxylamine to Fez+. For the sols with the smaller sized particles, several filtrations were required. Most experiments were conducted at pH 3.5, adjusted with HC104. For pH 6-7, phosphate buffen were used (2 X lo4 M). It was verified that the addition of the buffer had no effect on the yield of Fe(I1) from the reduction of the transparent sols.
Results and Discussion The reductant utilized in all experiments of the present study was the viologen radical V-. Its deep-blue color,A( = 602 nm; t 1.28 X lo4 M-' cm-l)17and reversible, pH-independent, one-electron reduction (EO = -0.41 V)18 make it a convenient electron-transfer mediator in radiolysis. The following sequence of reactions converts all the primary radicals produced by the radiolysis of aqueous solutions into V-: H20
--
e-(aq) (3.1); OH (3.1); H (0.6);H202 (0.7) (1)
e-(as)
+ N 2 0 + H20
OH(H) + (CHJ2CHOH (CH3)2COH+ V
-
-
-
OH
+ N2 + OH-
(2)
(CH3)ZCOH + H2O(H2) (3)
V-
+ (CH&CO + H+
(4)
The values in parentheses in eq 1 indicate the yields of the (14) Dimitrijevic, N. M.; Savic, D.; Micic, 0. I.; Nozik, A. J. J. Phys. Chem. 1984,88,4278. (15) Willner, I.; Ford, W. E. J . Heterocyclic. Chem. 1983, 20,1113. (16) Hart, E. J.; Anbar, M. The Hydrated Electron; Wiley: New York, 1969. (17) Willner, I.; Yang, J.-M.; Laane, C.; Otvos, J. W.; Calvin, M. J. Phys. Chem. 1981,85, 3277. (18) Nahor, G.S.;Rabani, J. Radiat. Phys. Chem. 1987, 29,79.
Mulvaney et al.
1208 Langmuir, Vol. 4, No. 5, 1988
2
3
4
5
6
7
8
PH
Figure 2. G(Fe2+)as a function of H for (A) 0.26 g/L a-Fe20,; (B) 0.26 g/L a-FeOOH; (C) 4 X 10-f M transparent a-Fez03sol. Total dose 10.2 krad, measured 0.5 h after irradiation. The solution composition was the same as that described in Figure 1.
respective species in NzO-saturated solutions in units of molecules/100 eV (denoted G values). The maximum yield of V- and thereby of reduction of Fe(II1) can thus be expected to be 6.8 molecules/100 eV. Reactions of hydrogen peroxide with either V- or with the Fe(I1) produced during the radiolysis are not expected to affect this yield. Both will regenerate OH in a Fenton-type reaction Fez+ (V-)
+ HzOz
-
Fe3+ (V)
+ OH + OH-
2
3
4
5
6
8
7
PH
Figure 3. Percentage of adsorbed Fe(I1) as a function of pH. Data taken from Figure 2.
N
(5)
followed by reactions 3 and 4. As described below, we did find G(Fe2+),, = 7.0 in the y-radiolysis experiments at low pH. A similar yield was obtained by Buxton et al.1° at pH 2.2. Reductive Dissolution of Iron(II1) Oxides by yRadiolysis. The yield of Fez+following y-radiolysis was measured for the colloids described in the Experimental Section under various conditions. Figure 1 shows that G(Fe2+)is independent of the initial concentration of the oxide in the range (0.2-2) X M. Since V- radicals are indefinitely stable in these solutions in the absence of the oxide, the observation of Figure 1 is not surprising. However, we do note in Figure 1that G(Fe2+)= 4.8 at pH 3.8 is substantially lower than G(Fe2+), = 6.8. Since no V- radicals were observed at the end of the irradiation period) three possibilities may be considered to explain the apparently low G(Fe2+)value. The first possibility is a surface-mediated side reaction that consumes V- but does not lead to reduction of Fe(II1). This was ruled out since both we and Dimitrijevic et al.14 found that although little Fe2+(aq)could be detected when the colloid was irradiated at relatively high pH (>8) quantitative amounts of Fe2+(aq) were recovered when the colloid was destroyed by acidification of the solution under an inert atmosphere. The other two possibilites are adsorption on the surface and trapping of Fe(1I) in the interior of the particles. To address the latter two possibilities some additional experiments were carried out and are described below. The yield of Fez+as a function of pH from reduction of cu-Fe203 and a-FeOOH suspensions and the small-sized particles of the transparent ol-Fez03sol is shown in Figure 2. As might be expected if adsorption of Fe(I1) causes the reduction in G(Fe2+),the yield indeed decreased with increasing pH. The same data are plotted in Figure 3 as the percentage of Fe(I1) remaining on the particles, assuming G(Fe2+)mar = 6.8. It can also be seen in Figures 2 and 3 that the amount of Fe(I1) remaining on the particle at the time of measurement is lowest for the transparent sol. This is not a result of differences in surface area. On the contrary, surface areas at the concentrations used in Figure 2 were 1.8, 4.8, and 14.7 m2/L for the a-FeOOH and a-
01 0
10
26
50 TIME (HOURS)
30
40
60
70
I
80
Figure 4. G(Fe2+)as a function of time after acidifying the irradiated Fez03 sol. pH at which the irradiation was carried out: (A) 6.9, (B) 10.7. Solid data points, solution acidified to pH 2.3; open data points, solution acidified to pH 3.5.
Fez03suspensions and the a-Fe203sol, respectively. As can be seen, the trend in the amount remaining on the particles is just the opposite. It seems, therefore, that the main reason for the low yields in the acidic range is not merely adsorption at the surface; charge migration into the particles' interior has to be considered. To further distinguish between the two possibilities, namely surface adsorption and deeper trapping) we irradiated a sol of cu-Fe203(4 X lo4 M) and a suspension of a-Fez03at pH 6.9 (0.26 g/L); portions of each of these colloids were then filtered and checked for Fez+. No Fez+ was found in either filtrate. The remainder of each of these irradiated colloidal solutions was then acidified under N2 to pH 2.2 and filtered again. G(Fe2+)= 5.2 was then found for the a-Fez03transparent sol, and G(Fe2+)= 2.0 was found for the a-Fez03suspension. While it is possible that Fe(I1) could partially be leached out by the acidification procedure) it is clear that some was trapped within the lattice structure. Direct irradiation at pH 2.2 did yield quantitative Fez+ formation (G(Fe2+),, = 7.0) for the transparent sol. It should also be emphasized here that any Fe(I1) in the bulk of the solution will predominantly exist as the Fe2+(aq)ion at any pH lower than 7 under all our experimental conditions.lg The rate of release of Fez+ from the particles was addressed in the following experiments. The transparent sol was irradiated at pH 6.9 or 10.7 with known doses. At the end of the irradiation, the pH was lowered (under N2 to prevent autooxidation),and the yield of Fez+was measured (19) (a) Baes, C. F.; Mesmer, R. E. The Hydrolysis of Cations; Wiley: New York, 1976. (b) Garrels, R. M.; Christ, C. L. Solutions, Minerals and Equilibria; Harper and Row: New York, 1965. (c) Encyclopedia of Electrochemistry of the Elements; Bard, A. J., Ed.; Marcel Dekker: New York, 1982; Vol. IX, Part A.
Reductive Dissolution of Iron(II0 Oxides
as a function of time. Results are shown in Figure 4. Acidification to pH 3.5 or 2.3 leads to the same yield of desorbable Fe(II), as can be seen in Figure 4. Except for a slight increase in the measured Fe2+yield at the early stage of the experiments at pH 3.5, no time dependence is observed. We therefore conclude that the rate of release of surface-adsorbed Fe(I1) upon acidification is fast on the time scale of these experiments, while the rate of release of Fe(I1) trapped in the bulk of the particles is very slow on this time scale. Irradiation of the solution at even more alkaline pH (10.7) leads to a decrease in the yield of desorbable Fe2+,which by inference means that the yield of charge trapping within the particles increases. Comparison between the yields of Fe2+upon irradiation at pH 2.2 vs irradiation at pH 6.9 or 10.7 followed by acidification to pH 2.2 leads to the conclusion that the pH of the solution is an important parameter in determining the fraction of charges trapped at the surface and the fraction that is able to escape such trapping. The fact that only a small yield of Fe2+could be recovered following the irradiation at the higher pH indicates that a large fraction of charge was able to migrate into the particles' interior. Thus, protonation at the surface promotes surface trapping. Since the rate of protonation will increase with acidity of the solution, lower pH will lead to an increase in yield of trapping at the surface. Of course, as is also clear from the adsorption isotherms described in the following section, trapping at the surface at high pHs does not lead to the release of Fe2+(aq). This is primarily because of the formation of ferrous hydroxides, which precipitate onto the surface. Thus the yield of Fe2+measured in the solution will always be higher at lower pH than at higher pH. Nevertheless, the experiments described above indicate that at higher pH charge migration into the bulk of the solid phase is more efficient than it is at low pH. The effect of pH may be amplified by the existence of the electric double layer at the oxide-H20 interface. A t pH values lower than the pzc, the positive surface potential would tend to enhance surface trapping while negative surface potentials will enhance charge migration into the bulk of the particles. Note that the pzc of the colloidal iron oxide used in the present study was measured to be at 8.3 f 0.3, in good agreement with literature values.20 In order to check the effect of surface potential on surface trapping we measured the effect of ionic strength on the yield of Fe2+. When a-FeOOH suspensions were irradiated at pH 3.0, the yield increased somewhat, frm G(Fe2+)= 2.1 in the presence of M NaC104 to G(Fe2+)= 2.5 in 1 M NaC104. For an amphoteric surface (such as the iron oxide surface), an increase in ionic strength leads to a decrease in surface potential.,l Concomitant with the reduction in surface potential will be an increase in the surface proton concentration following the Boltzmann distribution and an increase in the number of protonated surface sites (OH,'). Hence, the small increase in G(Fe2+) at the higher ionic strength suggests the surface potential is not the major factor in surface trapping. If the surface potential were the dominant factor governing trapping of electrons at the surface, then a decrease in G(Fe2+)would be expected. Another important parameter in determining the relative ratio between the two different trapping processes is the specific surface area of the particles. This is evident (20) (a) Parks, G.A. Chem. Reu. 1965, 177, 65. (b) Breeuwsma, A.; Lyklema, J. Discuss. Faraday SOC.1971,52, 324. (21) Chan, D. Y. C. In Chemical Processes At Mineral Interfaces; Davis, J. A,, Hayes, K. F., Eds.; ACS Symposium Series 323, American Chemical Society: Washington, DC, 1986, pp 99-112.
Langmuir, Vol. 4, No. 5, 1988 1209
;
E
20
0 E 40\ O
2
0
1
2
3
4
5 6 PH
7
8
9
1
0
Figure 5. Percentage of Fez+(total 4.4 X lo5 M Fez+)adsorbed on suspensions of a-FeOOH (0.26 g/L) and on transparent a-Fe03 M) as a function of pH. Circles are for a-FeOOH, sol (1 X and squares are for a-FeZO3.
on comparing the results from the two colloids under the same experimental conditions. For the smaller sized particles with their higher surface/volume ratio, surface trapping invariably exceeds that of the suspensions. When irradiated at near neutral pH, our results indicate that ca. 25% of the transferred electrons is trapped within the crystal lattice for the transparent sol. For the suspensions, ca. 70% of the charge transferred migrates into the bulk of the colloid. Chemical identification of the trapping sites cannot be verified by the experiments performed in this study. We may, however, note that Kennedy and Frese22 identified two donor states in polycrystalline a-Fe20g. One state is located very near the flatband potential and a deeper state, which is 0.6 V positive of the flat band potential. It is conceivable that these two states are associated with the two trapping sites discussed above. Adsorption of Fe(I1). Adsorption of Fe(I1) ions onto the iron(II1) oxide colloidal solutions was studied under various oxide and Fe2+concentrations and at various pHs. Solutions similar to those used in the irradiation experiments were used in this series of experiments, and the time of determination of free Fe2+remaining in the solution was even longer than that in the steady-state radiolysis experiments. Figure 5 shows the results from these adsorption experiments for transparent a-Fe203sols and a-FeOOH suspensions at various pHs. A comparison between Figures 5 and 3 reveals that the fraction of Fe2+ adsorbed in the experiments of Figure 5 is much smaller than the fraction in Figure 3. Even for the very small particles of a-Fe203, only 5% of the total Fe2+added to the sol adsorbs on the particles (Figure 5) at pH 4, while more than 25% remains on the irradiated sol of Figure 3. Note that the concentration of cu-Fe203in the adsorption experiments was higher than that in the irradiation experiments. For the larger particles of the suspensions used, the differences are even more pronounced. Superposition of Figure 5 on Figure 3 thus reveals a pronounced hysteresis between the release experiments of the irradiated colloids and the adsorption isotherms. Clearly, thermodynamic equilibrium was not achieved in at least one of these experiments. Although the two types of experiments showed total adsorption at high pH, this results primarily from precipitation of iron(I1) hydroxides on the particles' surface in the adsorption experiments. In the irradiated solutions, we have already shown above that most of the Fe(I1) is not at the surface. While this hysteresis may result from either slower adsorption or from slower release of Fe2+we believe the latter to be the case. The fact that little further Fe2+is released more than an hour after the acidification in Figure 4 precludes the likelihood that desorption hysteresis is due to a gel layer at the surface (22) Kennedy, J. H.; Frese, K. W. J.Electrochem. SOC.1978,125,723.
1210 Langmuir, Vol. 4 , No. 5, 1988
Mulvaney et al.
Table I. Thermodynamic Data Used in Construction of the Eh-pH Diagram (Figure 6)19 thermodynamic parameter
reaction
+
FeZt(aq) + HzO = FeOH+ H+ Fe2+(aq)+ 2 H z 0 = Fe(OH), + 2Ht Fe3+(aq)+ HzO = FeOHZt + H+ 2Fe3+(aq) 3 H z 0 = a-FezO3 6H+ Fe3+(aq)+ 2 H z 0 = am-FeOOH + 3Ht Fe203 6HC 2e- = 2Fe2+ 3 H z 0 Fe304 8H+ + 2e- = 3Fe2+ + 4 H 2 0 Fe3+ e- = Fez+ Fe2+ + 2e' = Fe Fe(OH), 2H+ 2e- = Fe 2 H 2 0 3FezO3 2H+ 2e- = 2Fe304 HzO Fe304 + 2H20 2H+ 2e- = 3Fe(OH),
+
+ + +
+ +
+ +
+
+ + +
+
+
+
log Kl1 = -9.5 log K12 = -12.85 log K,, = -3.05 log Kslo = 1.88 log Kslo = 3.55 Eo = +0.728 V Eo = +0.980 V Eo = +0.771 V Eo = -0.44 IF' = -0.087 V Eo = +0.221 V E" = -0.197 V
v
within which the Fe2+is distributed. Adsorption of metal ions on ferric oxide colloids has been extensively s t ~ d i e d . ~In~addition , ~ ~ to the possibility of bidentate bonding and simultaneously adsorption and hydrolysis, adsorption of Fe(I1) may lead to slow formation of magnetite26at high pH. However, the migration of Fe(I1) into the lattice of the oxide is even slower than the time scale involved in the irradiation experiments. The mechanism for trapping of Fe(I1) states in the bulk of the solid following irradiation is therefore different from the one in the adsorption experiments. Thermodynamics of Iron(II1) Oxide Reduction. In the preceding sections we conclude that viologen radicals react with colloidal iron(II1) oxides to produce the three possible Fe(I1) species: Fe2+(aq)in solution, Fe(I1) adsorbed at the particle's surface, and Fe(I1) trapped in the particle's interior. Only at high acidities (i.e., pH C3) and large surface/volume areas does the concentration of Fe2+(aq)approach the value predicted by the stoichiometry of eq 6. We now analyze the thermodynamics of the Fe(III)/Fe(II) system through detailed examination of the Pourbaix diagram constructed from known thermodynamic data and compare it with our results. The Eh-pH diagram for the relevant systems is shown in Figure 6. In addition to the various Fe(III)/Fe(II) species, the V/V-, H+/H2, and 0 2 / H 2 0couples are also plotted in this figure. Data used to construct this diagram are provided in Table I. The small Fe3+field compared with the wide Fe2+field in this diagram provides the rationale for the use of reductive rather than protonic dissolution of the oxides. The solubility of the oxide increases by up to lo8 when dissolution is into the ferrous form,4c,26 and much milder conditions are needed to dissolve the oxide. Of course, formation of the Fe2+center leads to an increase in ionic radius and changes in coordination that also serve to accelerate the expulsion of the ferrous ion from the l a t t i ~ e .Reductive ~,~ dissolution is consequently favored both thermodynamically and kinetically. From Figure 6, it can be seen that at any pH 110.5 V- is thermodynamically capable of reducing Fe203. Indeed, we observe complete disappearance of V- at pH 57, while at pH >9.5 the blue coloration due to V- formation in the irradiated sols persists for minutes. At pH 11.2 we find almost negligible loss of viologen radicals produced by radiolysis. (23) Benjamin, M. M.; Leckie, J. 0. J. Colloid Interface Sei. 1981, 79, 209. (24) (a) Hohl, H.; Stumm, W. J. Colloid Interface Sei. 1976,55, 281. (b) Loganathan, P.; Burau, R. Geochim. Cosmochim. Acta 1973,37,1277. (25) Tamaura, Y.; Ito, K.; Kataura, T. J. Chem. Soc., Dalton Trans. 1983, 189. (26) Yariv, S.; Cross, H. Geochemistry of Colloid Systems; SpringerVerlag: Berlin, 1979.
1 H 2 0 / 0 2 dabilily
IRON - WATER SYSTEM @ 298K
2 H20/H2 stability
10
A Fe203 E Fe2+ 0.
.
..
Eo N/V
.0.5
-)
.1.0 0
2
4
6
8
10
12
14
16
PH
Figure 6. Eh-pH diagram of the Fe(II)/Fe(III), V / V , H+/H2, and H 2 0 / 0 2systems constructed with the parameters listed in Table I.
Under typical experimental conditions of the radiolysis experiments, direct reduction to Fe2+(aq)according to eq 6 is possible only at pH
