Langmuir 1991, 7, 1627-1634
1627
Light-Induced Dissolution of Hematite in the Presence of Oxalate: A Case Study Christophe Siffert and Barbara Sulzberger’ Institute for Water Resources and Water Pollution Control (EAWAG), Zurich, Swiss Federal Institute of Technology (ETH), Zurich, Switzerland Received November 19,1990. In Final Form: February 18, 1991 The light-induceddissolutionof hematite in the presence of oxalateoccurs (a)through the photochemical reductive dissolution of hematite resulting in formation of dissolved Fe(I1) and oxidized oxalate and (b) through the Fe(I1)-catalyzedthermal dissolution of hematite, resulting in formation of disaolved Fe(II1). These two dissolution pathways are coupled via the photochemical reduction of dissolved iron(III), i.e. photolysis of iron(II1) oxalato complexes, leading to an autocatalytic dissolution of hematite. The rate of the photochemicalreductive diasolutionof hematite is strongly wavelength-dependent,and the reaction occurs only in the near-UV (A < 400 nm). From this wavelength-dependencewe conclude that the electronicallyexcited state involved in this heterogeneousphotoredox reaction is a ligand-to-metalchargetransfer state, either of the iron(II1) oxalato surface complex or of the bulk hematite (or both). Oxygen strongly inhibita the reductive dissolution of hematite under the influence of light. This phenomenon is interpreted in terms of competition between reoxidation of reduced surface iron by oxygen and phase transfer of surface Fe(I1) into solution. These observations indicate that in this system detachment of surface Fe(I1)from the crystal lattice is the rate-determiningstep of the overall dissolution process rather than electron transfer. In the presence of oxygen hematite acts as a photocatalyst for the oxidation of oxalate by oxygen. The rate of this photocatalytic oxalate oxidation exhibits a similar wavelengthdependence as does the photochemical reductive dissolution of hematite.
Introduction The kinetics of the dissolution of iron(II1) (hydr)oxides playa an important role in the cycling of iron in natural waters, which is coupled to the biogeochemical cycling of many other chemical compounds such as heavy metals and nutrients.’ Reduction of solid phase Fe(II1) greatly enhances the dissolution because of the larger lability of the F e w crystalline lattice bond as compared to the F e u crystalline lattice bond. Reductive dissolution of iron(II1) (hydrloxides occurs primarily at the sedimentwater interface under anoxic conditions in the presence of reductants, such as products of the decomposition of biological material or exudates of organisms. Reductive dissolution of iron(II1) (hydr)oxides, however, can also occur in the photic zone and in atmospheric water in the presence of compounds that are metastable with regard to iron(III), that is, compounds that do not undergo redox reactions with iron(II1)unless catalyzedby light. Dissolved oxygen is usually the oxidant of iron(I1). If the lightinduced dissolution of iron(II1) (hydrloxides occurs with a sufficiently high efficiency,a steady-state concentration of dissolved iron(II) may be maintained. Diurnal variation in the Fe(I1) concentration has been observed in acidic surface The photoredoxcyclingof iron increases its bioavailability to the phytoplankton: The oxidative hydrolysis of the photochemically formed iron(I1) yields oxidation products that are much more readily thermally dissolved than are the crystalline iron(II1) (hydr)oxide phases, resultingin formation of dissolved iron(II1) species, which are thought to be the important species for the uptake of iron by the biota.6 The formation of bioavailable iron species is of particular significancein areas where primary productivity is iron-limited.6 In natural waters (1)Zinder, B.;Stumm, W. Chimia 1986,39,280. (2)Colienne, R.H. Limnol. Oceanogr. 1983,28,83. ( 3 ) McKnight, D. M.; Ben&, K. E. Arctic Alpine Res. 1988,M,492. (4)Sulebever,B.;Schnoor, J. L.; Giovanoli, R.; Hering, J. G.; Zobriet, J. A p w t . SCL1990,62,M. (5) Rich, H. W.; Morel, F. M. M. Limnol. Oceonogr. 1990,35,349. (6) Martin, J. H.; Fitzwakr, S. E. Nature 1988,331,341.
various redox pairs are often not in equilibrium with each other, e.g. 0 2 in the presence of HSOs- or organic solutes. Thus, the photoredox cycling of iron in surface waters may enhance the oxidative degradation of organic compounds, either biogenic or xenobiotic, and hence also be of importance for the abiotic degradation of organic pollutants. In this paper we will present a case study on the kinetics of the light-induced dissolution of iron(II1) (hydr)oxides with hematite as the solid phase and oxalate as the electron donor. We will illustrate (1) that coupled reaction pathways are involved in the light-induced dissolution of hematite leading to an autocatalytic dissolution process, (2) that a ligand-to-metal charge-transfer state is involved in the photochemical reductive dissolution of hematite in the presence of oxalate, and (3) that detachment of the reduced surface iron is likely to be the rate-determining step of the overall dissolution process.
Background In aquatic systemsthe photochemical iron(I1)formation can occur via different pathways7** shown in Figure 1: (1) through photochemical reductive dissolution of iron(II1) (hydr)oxides;(2) through photolysis of dissolved iron(II1) coordination compounds.10 Dissolved iron(II1) (i) is an intermediate of the oxidative hydrolysis of Fe(I1) and (ii) results from the thermal nonreductive dissolution of iron(111) (hydr)oxides, a reaction that is catalyzed by iron(11).11J2 The various hypothetical steps involved in the photochemical reductive dissolution of an iron(II1) (hydrloxide (7)Waite, T.D.;Morel, F. M. M. J. Colloid Interface Sci. 1984,102, 121. (8)Waite, T.D.; Morel, F. M.M. Enuiron. Sci. Technol. 1984,18,880. W.; Mor an, J. J. Apwrtic Chemistry: An Introduction (9)St-. EmphusiztngChemical ~puilibriain Natural Waters:Wiley-Intsrsdence: New York, 1981. (10)Balzani, V.; Carassiti, V. Photochemistry of Coordination Compounde; Academic hew: London, 1970. (11)Sukr, D.;Siffert, C.; Sulzberger, B.;Stumm, W. Natunuissenschaften 1988,76, 671. (12)Suter, D.;Banwart, S.; Stu”, W.Langmuir 1991,7,809.
0743-746319112407-1627%02.6010 0 1991 American Chemical Societv
1628 Langmuir, Vol. 7, No. 8, 1991
Siffert and Sulzberger
ligand I-
Figure 1. Schematic representation of the aquatic photoredox cyclingofiron.g>denoteethelatticesurfaceofaniron(II1)(hydrloxide, and ET stands for electron transfer.
tl
k2
I
H.15
Surlace complex formation
Diesodation and decarboxylation of the oxidizedoxalate
Detachment of surface Fe(li)
involved in this heterogeneousphotoredoxreaction. After dissociationfrom the surface,the oxalate radical undergoes a fast decarboxylation reaction's yielding COZand the C0z'- radical, which is a strong reductant and can reduce a second surface iron(II1) in a thermal reaction. For the sake of simplicity,this thermal reaction of the CO$-radical is omitted in Figure 2. We assume that detachment of the reduced surface iron occurs as a last elementary step of this dissolutionprocess. If the surface complex is the chromophore, then the photochemical reductive dissolution occurs as an intramolecular process (mechanism a); alternatively, if the bulk iron(II1) (hydr)oxideis the chromophore, then it is an intermolecular process (mechanism b). The elementary steps and the corresponding rate expressionsfor mechanismsa and b are given in parts a and b of Table I, respectively. From the rate expressions it can be seen that irrespective of whether the surface complex or the bulk iron(II1) (hydr)oxideacta as the chromophore,the rate of dissolvediron(I1)formation is directly proportional to the surface concentration of the specifically adsorbed electron donor. It has been shown experimentally with various electron donors that the rate of dissolved iron(I1) formation under the influence of light is a Langmuir-typefunction of the dissolved electrondonor concentration.me2' Photolyses of dissolved iron(II1)oxalato complexes are well-knownreactions'o that occur with high quantum yields according to the following overall stoichiometry for the 1:l complex in the absence of oxygen:
hu
2 Fe"C20,+ 2Fe2++ 2C0, + CO , -: (13) The rate of this homogeneous photoredox reaction is directly proportional to the concentration of the iron(II1) oxalato complex
R = k[FemC,O,+]
(14)
where Figure2. Schematicrepresentationof the various step involved in the photochemicalreductive dissolution of an iron(II1)(hydr)oxide in the presence of a ligand such as oxalate.
in the presence of a ligand such as oxalate are outlined in Figure 2. An important step is the formation of a surface complex. Inner-sphere surface coordination equilibria have been described in terms of surface ligand exchange r e a c t i ~ n sand ~ J the ~ ~ constant ~ capacitancemodel.g Spectroscopic investigationswith magnetic resonance methods (EPR, ENDOR, ESEEM),lBwith EXAFS,'' and with FTIR18 have confirmed the inner-sphere structure of different surface complexes. Light absorption results in an electronicallyexcited state (indicated with an asterisk). Electron transfer via this electronicallyexcited state leads to a reduced surface iron center and the oxidized oxalate. It is likely, but not necessarily, the surface complex that acta as the chromophore (as suggested in Figure 2); it cannot be excluded, however, that an electronicallyexcited state of the bulk iron(II1) (hydr)oxide (or both) might be (13) Stu", W.; Kummert, R.; Sigg, L. Croat. Chem. Acta 1980,53, 291. (14) Schindler, P. W.; Stumm, W. In Aquatic Surface Chemistry; Stumm, W., Ed.; Wiley-Inbmience: New York, 1987; Chapter 4. (15) Dzombak, D.; Morel, F. M. M. Aquatic Sorption: Stability Constantsfor Hydroud Ferric Oxido;Wiley-Inbdence: New York, 1990. (16) Mobchi,H.In Aquatic Surface Chemistry;Stumm,W.,Ed.;WileyInbmience: New York, 1987; Chapter 6. (17) Hayes, K.F.; Roe, A.L.;Brown, 0.E.,Jr.; Hodgeon,K.0.;Leckie, J. 0.;Parka, G. A. Science 1987,298,783. (18) Zeltner, W. A,; Yoet, E. C.; Machesky,M. L.; Tejedor-Tejedor,I.; Andemn, M. A. In Ceochemrcal Processes at Mineral Surfaces; Davis, J. A., Hayes, K.F., E%.; American Chemical Society: Washington, DC, 1986.
k = 2.3L10,c,@, for 2.3Lc,[FemC,0,+] FemOH surface complex >Fe"'OH
E >Fe(II) + OH
(17) This surface photoredox reaction is analogous to photolysis of the dissolved FeOH2+complex
Langmuir, Vol. 7, No. 8, 1991 1631
Dissolution of Hematite in the Presence of Oxalate h
1.a
"
h
5E
f l '
.e
0
0.8
E
W
W
4.
0.6
I
-c8
*J 0.4
g
0.2
0.0
3
0
6
9
time
12
L-1.
6
4
(h)
time
Figure 6. Light-induceddiesolutionof hematite in the presence of oxalate at pH 3. The deaerated hematite suspension was irradiated with light that had passed the monochromator (X = 375 nm;ZO = 4 W/m2). Initial oxalate concentration was 3.3 "01
2
0
(h)
Figure 6. Light-induceddissolution of hematite and photooxidation of oxalate in a deaerated hematite suspension at pH 3 white light, ZO= 4 kW m4; initial oxalate concentration,1mmol L-1.
hr
FeOH2+ Fe2++ 6 H (18) The solution species FeOH2+exhibits a ligand-to-metal charge-transfer band with a maximum at 295 nm.90 The quantum yield of reaction 18has been reported to be 0.14 at 313 nm and 0.017 at 360 nmem On the other hand the quantum yield of photolysis of ferrioxalate complexes is 1.2 at 365 nm.28 It is likely that there is alsoa big difference in the quantum yields of photolysis of the corresponding surface species in the near-UV, which would be one reason for the enormous effect of oxalate on the reductive dissolution of hematite upon irradiation with near-UV light. Figure 5 shows the effect of the superposition of different pathways on the kinetics of the light-induced dissolution of hematite in the presence of oxalate. The experiment shown in this figure was carried out at a relatively low light intensity. Under these experimental conditions not only dissolved iron(I1) is formed but also dissolved iron(111). The rate of dissolved iron(II1) formation is constant within the time scale of this experiment; the rate of dissolved iron(I1) formation, however, increases with time. The appearance of dissolved iron(I1) and the disappearance of oxalate upon irradiation of a deaerated and an aerated hematite suspension are shown in Figures 6 and 7, respectively. In adeaerated suspension the reaction occurs according to the following overall stoichiometry:
+
a-Fe20s C,O,'
hr
+ 6H+ N2 2Fe2++ 2C0, + 3H,O
(19)
This is no longer true in the presence of oxygen, as shown in Figure 7. In this case, oxalate is oxidized much faster than under nitrogen and the concentration of dissolved iron does not exceed 0.04 mmol L-l at the maximum. At pH 3 and in the presence of oxalate, the solubility of iron(111) is much higher than 0.04 mmol L-1.O Furthermore, the ferrioxalate species are readily photolyzed. Thus, the lack of dissolved iron cannot be explained in terms of oxidative hydrolysis of Fe(II),but rather by the inhibition of the reductive dissolution. Plausibly, adsorbed molecular oxygen reoxidizes reduced surface iron
-
>Fe(II) + 0, 02*+ >Fe(III) (20) Thus, in the presence of oxygen, hematite acts as a photocatalyst for the oxidation of oxalate
-
"
1
2
[Feldissolved 0.0 0
-
I
3
4
. 5
time (h)
Figure 7. Photocatalytic oxidation of oxalate in an aerated hematite suspension at pH 3 white light, ZOz 4 kW m-2; initial oxalate concentration, 1mmol L-l.
CO , -:
+ 2H' + 0,
hu
heautib
H,O,
+ 2C0,
(21)
Oxidation of oxalate with 6 H radicals (formed through the Fenton reaction) in addition to its heterogeneous photochemical oxidation, and oxidation of the intermediate species, C02'-, by oxygen may account for the higher rate of oxalate disappearance in an oxygen-saturated suspension as compared to the rate in a deaerated suspension. Figure 8 shows the wavelength dependence of the rate of photochemicalreductive dissolution of hematite in the presence of oxalate at constant incident light intensity. As seen in Figure 8 only light in the near-UV region (A < 400 nm) leads to an enhancementof the dissolution brought about by the redox process at the surface. From these rates we calculated the quantum yields for mechanism a and b according to the eq 4 and 11,respectively. These quantum yields are listed in Table 11. For the calculation of the quantum yields, based on the assumption that the Fe(II1) oxalato surface complex is the chromophore, we used the extinction coefficients as estimated from the absorption spectrum of dissolved ferrioxalate. Because of the muchlower concentration and presumably the much lower extinction coefficients of the iron(II1)oxalatosurface
Siffert and Sulzberger
1632 Langmuir, Vol. 7, No. 8, 1991
0 350
360
370
380
wavelength
390
400
1
(nm)
Figure 8. Rate of the hotochemical reductive dissolution. of hematite, Rx dFe(II)/&,in the presenceof oxalate aa a function of the wavelength at constant incident light intensity ( l o = loo0 peinsteins L-l h-9. The hematite suspensions were deaerated; initial oxalate concentration = 3.3 mmol L-l; pH = 3. (In order to keep the rate of the thermal dissolutionconstant,a high enough concentration of iron(II),[Fez+]= 0.15 mmol L-l, was added to the suspensions from the beginning. Thus, the rates correspond to dissolution rates due to the surface photoredox process.)
Table 11. Quantum yields, Ob of Dissolved Iron(I1) Formation through the Heterogeneous Photoredox Reaction.
350
360
370
380
390
wavelength
400
410
420
430
(nm)
Figure 9. Rate of the photocatalytic oxidation of oxalate, RA= -d[C2OP]/dt, in aerated hematite suspensions aa a function of the wavelength at constant incident light intensity ( l o = loo0 peinsteins L-l h-1). Initial oxalate concentration = 1mmol L-l; pH z 3. (Note: In these experiments, only a small percentage of the initial oxalate was oxidized. Because of the uncertainty associated with the determination of the extent of oxalate oxidation, the rates are only accurate to about &20% .)
a-Fe203+ C20,2- + 6H+ e 2Fe2++ 2C02 + 3H20 (22)
The abovevalue for AGO is based on the followingsolubility product of hematite: pFe 3pOH = 42.7.= This is unlike lox, RA= @Ab = reductive dissolution of manganese(II1,IV) (hydr)oxides A, reinsteins d[Fe*+]/dt, 01. = R,/h nm L-l h-l rmol L-' h-1 R A / ~ . ~ L Z O ~ ~ A ~ > F ~ I(10-9) I I C ~ O , - } in the presence of oxalate, which occurs with high rates in the dark.a*%These differences might be explained in 350 766.8 27.0 1.4 35.0 terms of the higher thermodynamic driving force in the 360 986.4 29.1 1.5 29.5 case of Mn(II1,IV) (hydr)oxides and by the differences in 375 1090.8 17.0 1.4 15.6 the electronic structure of Fe(II1) and Mn(III).s*S8The 390 1321.2 12.2 1.5 9.2 405 1292.4 0.9 0.2 0.7 effect of light in the hematite/oxalate system is of catalytic nature: Formation of an electronically excited state a The quantum yields @A* are calculated under the assumption enables the system to overcome the free energy of that thesurface complex is the chromophore: {>F&204-} = 2Opmol L-1; L = 0.5 cm; the cA values are estimated from the absorption activation of electron transfer (see Figure 10). spectrum of dissolved ferriodate, and hence these quantum yields The wavelength-dependence of the rate of photochemcan only be considered as very rough estimations. The quantum ical reductive dissolution of hematite was studied in order yields are calculated under the assumption that the bulk iroto get information about the chromophore(s) and the n(II1) (hydr)oxideis the chromophore. electronically excited state(s) involved in this heterogeneous photoredoxprocess. The quantum yields as defined complex as compared to the hematite bulk we by eqs 5 and 12 represent the probabilities that a phowere not able to measure the absorption spectrum of the toredox reaction from an electronically excited state takes surface complex, since it is masked by that of hematite. place, leading to surface iron(I1) and the oxidized oxalate. The wavelength dependence of the rate of photocataThe quantum yields, estimated on the basis of eq 4 are lytic oxidation of oxalate a t constant incident light much higher than those calculated according to eq 11.To intensity in aerated hematite suspensions is shown in conclude from the differences between P and @b, however, Figure 9. The rate is within 20% constant between 430 that the surface complex is mainly involved as chroand 390 nm and increases below 390 nm. The fact that mophore, would be a vicious circle, since the quantum the rate does not drop to zero at wavelengths above 400 yields, calculated according to eq 4 are bound to be much nm, as is the case for the photochemical reductive higher than those calculated according to eq 11under our dissolution of hematite in deaerated suspensions, may be experimental conditions. If the surface complex were the explained in terms of photolysis of dissolved iron(II1) oxresponsible chromophore, then one would expect the alato complexes that are formed through thermal dissoquantum yields to be constant in the wavelength range of lution of hematite and are photolyzed with high quantum its ligand-to-metal charge-transfer band as are the quanyields up to 500 11111.28 tum yields of photolysis of dissolved iron(II1) oxalato complexes for an individual iron(II1) oxalato species.29 Discussion Between 350 and 390 nm the estimated quantum yields Reductive dissolution of hematite in the presence of oxalate occurs as a photochemical process and not as a (33) Stucki,J. W.; Goodman, B. A.; Schwertmann, U. Iron in Solids thermal reaction, despite the relatively high thermodyand Clay Minerale; D. Reidel: Dodrecht, 1988. (34) Stone, A. T. Ceochim. Cosmochim. Acta 1987,61,919. namic driving force, AGO = -0.87 eV at pH 3 for the (35) Xyla,G. A.; Sulzberger,B.; Luther, G.W., III; Hering, J. G.; VM following equilibrium: (32) Marusak, L. A,; Meeeier, R.;White, W. B. J. Phys. Chem. Solids 1980,42,981.
+
Cappellen, P.; Stu", W. Submitted for publication in Langmuir. (36) Luther, 111,G.W. In Aquatic ChemicalKinetics: ReactionRotes of h c e s s e s in Natural Waters; Stu", w., Ed.; Wiey-Interscience: New York, 1990; Chapter 6.
Dissolution of Hematite in the Presence of Oxalate
reaction coordinatei
Figure 10. Qualitative representation of the energetice of the photochemicalreductive dissolution of hematite in the presence of oxalate. > F a x is the iron(II1)oxalatosurface complex, i.e. the precursor complex, in ita electronically ground state and >FeOx*the precursor complex in ita electronicallyexcited state. AGO is the free energyof the overall dissolutionprocess according to eq 22; AC*m is the free energy of activation of formation of a reduced surface iron, >Fe(II), and the oxidized oxalate, and A C * Dis~ the free energy of activation of the detachment of the reduced surface iron from the crystal lattice. For the sake of simplicity, the oxidized product is omitted in this figure.
are practically constant. At 405 nm, however, cPa is a factor of 7 smaller than at 350 nm. This is unlike the wavelengthdependence of CP of photolysis of dissolved iron(II1) oxalato complexes, where the quantum yield at 405 nm is only 4% smaller than that at 350 nm.= The comparatively low value of @a at 405 nm is somewhat surprising, since one would expect the iron(II1) oxalato surface complex to exhibit an absorption spectrum that is rather red- than blue-shifted as compared to the absorption spectrum of dissolved f e r r i ~ x l a t e .If~ ~the bulk hematite phase were mainly involved as the chromophore in this heterogeneous photoredox reaction, then one would expect the quantum yields not to be constant over the wavelength range of the absorption spectrum of hematite: Hematite is a semiconductor with different bands and thus energetically forbidden zones of different energies.38 Electronic excitation of the 2.3-eV band gap corresponds to a ligand-field transition, whereas electronic excitation of the 3.3-eV band gap corresponds to a Fem O-n charge-transfer transition, i.e. a ligand-to-metalcharge-transfer transition, the lattice iron(II1) being the metal and the lattice oxygen being the ligand.32139 The absorption edge of the 2.3-eV band gap occurs at 540 nm and that of the 3.3-eV band gap at 375 nm. Since only light with X < 400 nm leads to a photochemical reductive dissolutionof hematite, electronic excitation of the 2.3-eV band gap does not appear to be operative as an oscillator in this heterogeneous photoredox reaction. From our data we conclude that either a FelI1 O-II charge-transfer transition of hematite or a ligand-to-metal charge-transfer transition of the surface complex (or both) are the oscillators that drive the redox reaction leading to reductive dissolution of the solid phase. These conclusions are consistent with those of Faust and Hoffmanna and Litter and Blesa" who investigated the wavelength-dependence of the rate of photochemical reductive dissolution of iron(II1) (hydr)oxides using hematite-bisulfite and maghemite-EDTA as model systems, +
-
(37) Regazzoni, A. E.; Blew, M. A. Lcrngmuir 1991, 7,473. (38) Coodenough,J. B. In Progress in Solid-State Chemistry; Reisa, H., Ed.;Pergamon: Oxford, 1971; Chapter 4. (39) Vaughan, D. J.; Toseell, J. A. Can. Mineral. 1978,16,159. (40) Fauet, B. C.; Hoffman, M. R. Enoiron. Sci. Technol. 1986,20, 943. (41) Litter, M. I.; B l w , M.A. J . Colloid Interface Sci. 1988,125,679.
Langmuir, Vol. 7, No. 8, 1991 1633 respectively. An unequivocal identification of the chromophore(s) and the electronicallyexcited state(s)involved in this heterogeneous photoredox reaction would require a suitable ligand that forms an inner-sphere surface complex with surfaceiron(III),exhibiting a ligand-to-metal charge-transfer band in a spectral window of the solid phase and being readi1y.photolyzed. The tremendous effect of oxalate on the photochemical reductive dissolution of hematite may be interpreted in terms of a much more efficient charge transfer to the surface bound oxalate than to the surface hydroxo goup. The efficiency of electron transfer in heterogeneous photoredox reactions depends strongly on the type of surface complex present. Using methylviologen (MV2+)as electron acceptor, Darwent and L e ~ r e 'have ~ reported that the yield of photochemicalMV'+formation with Ti02 as photocatalyst was much higher in the presence of an electron donor where two functional groups are likely to be bound to the Ti02 surface as compared to an electron donor that can only form a monodentate surface complex. Oxalate is likely to form a bidentate surface complex at the surface of hematite. A further reason for oxalate being such an efficient electron donor in the photochemical reductive dissolution of iron(II1) (hydr)oxides is that the oxalate radical undergoes a fast decarboxylationreaction,l8 such that back electron transfer becomes negligible. For the same reason, EDTA is also a very efficient electron donor and is often used as a sacrificial compound in photoredox reactions for photochemical solar energy conversi0n.4~The effect of oxalate on the steady-state concentration of dissolved iron(I1) has been demonstrated in an acidic lake with a low content in dissolved organic carbon (DOC 0.5 mg L-1).4 The rate of dissolved iron(I1) formation depends not only on the efficiency of electron transfer but also on the efficiency of detachment of reduced surface iron from the crystal lattice. In the presence of oxygen the photochemical reductive dissolution of hematite is inhibited. This phenomenon can be interpreted in terms of reoxidation of reduced surface iron by oxygen. With lepidocrocite as the solid phase, the rates of the photochemicalreductive dissolution in a deaerated and in an aerated suspension are almost identical at pH 3.u Lepidocrocite is an iron(111) (hydr)oxide phase that is thermodynamically less stable than hematite.= Thus, reduced surface iron ions are more readily detached from the lepidocrocite crystal lattice than from the hematite crystal lattice. These experimental results provide circumstantial evidence that in this system detachment of the reduced surface iron centers, which is in competition with their reoxidation by oxygen, is the rate-determining step of the overall photochemical dissolution process. The following mechanisms7can be invoked for the photocatalytic oxidation of oxalate in an aerated hematite suspension: (a) ligand-to-metal charge-transfer process within the iron(II1) oxalato surface complex, resulting in the oxidation of oxalate and the reduction of surface iron that is subsequentlyreoxidized by oxygen; (b)photo-Kolbe process& in which photo-generated holes are scavenged by oxalate and the photoelectronsare scavenged by surface iron that is subsequently reoxidized by oxygen; (c) oxidation of oxalate with OH radicals, resulting from the
=
(42) Darwent, J. R.; Lepre, A. J. Chem. SOC.,Faraday Tram. 2 1986, 82, 2323. (43) Tazuke, S.; Kitamura, N.; Kim, H. B. In Supramolecular Photochemistry; Balzani,V., Ed.;NATO AS1 Series, Ser. C, Vol. 214; D. Reidel Dordrecht, 1987. (44) Sulzberger, B.;Laubecher, H. U. In preparation. (45) Kraeutler, B.;Bard, A. J. J. Am. Chem. SOC.1978,100, 5986.
Siffert and Sulzberger
1634 Langmuir, Vol. 7, No. 8, 1991 photooxidation of surface OH gr0upe.M The third mechanism is unlikely under our experimental conditions, because of the low efficiency of photooxidation of surface OH groups in the near-W region. From the wavelengthdependence of the rate of the photocatalytic oxidation of oxalate at constant incident light intensity in aeratad hematite suspensions,we conclude that either mechanism a or b is involved in the photocatalytic oxidation of oxalate. In the case of a photo-Kolbe process, the photoholes are not formed through excitation of the 2.3-eV band gap of hematite but through excitation of the 3.3-eV band gap, corresponding to a LMCT (Fern 0") transition. The autocatalyticiron(II) formation4' as shown in Figure 5 can be explained with the following scheme:
-
t w2-
Dissolved iron(I1) is formed through the heterogeneous and the homogeneous photoredox reaction. The rate of dissolved iron(I1) formation is thus the s u m of the rates of Fe(I1) formation through the heterogeneous and the homogeneous photoredox reactions. The rate of the heterogeneousphotoredox reaction is directly proportional to the surface concentration of the specifically adsorbed oxalate, (>FemC20d-)! and the rate of the homogeneous photoredox reaction is directly proportional to the concentration of dissolved iron(II1) oxalato complexes, [Fem(C204),*]; compare eqs 4, 11, and 14. Under our experimental conditions the surface concentration of the specifically adsorbed oxalate can be aasumed to be constant. What increases with time, however, is the concentration of dissolved iron(II1) oxalato complexes and thus the rate of dissolved Fe(I1) formation. The rate of dissolved iron(II1) formation is equal to the rate of the Fe(I1)-catalyzed dissolution of hematite minus the rate of homogeneous photolysis of iron(II1) oxalato complexes. The rate of the Fe(I1)-catalyzed dissolution depends-at a constant oxalate concentration-on the surface concentration of the adsorbed Fe(II) through a bridging ligand; compare eq 16. Both rates increase with time, because of the increase in {Fe(II)jd and [Fem(C204)ss-]. But the difference in the rates is constant within the time scale of (46)Herrmann, J. M.;Mozzanega, M.N.;Pichat, P.J. Photochem.
--.
I om.--, 22.333. (47) ComeU, R.M.;Schindler, P.W.Clays Clay Miner. 1987,36,347.
this experiment. The rate of dissolved iron(I1) formation is expected to become constant as soon asthe concentration of dissolved iron(II1) oxalato complexes reaches a steady State
This is the case if the rate of Fe(C204)s9-formationthrough the Fe(I1)-catalyzed dissolution of hematite is equal to
Conclusions Rsductive dissolution of hematite in the presence of oxalate occurs as a photocatalytic process, i.e. the light energy is used to overcome the activationenergy of electron transfer. Depending on the light intensity and ita attenuation by the hematite particles, not only the surface photoredox reaction but also photolysis of dissolved iron(II1) oxalato species (that are formed through the thermal dissolution of hematite) takes place. Both photoredox reactions result ultimately in formation of dissolved iron(11)and CO2. Since Fe(I1) acta as a catalyst in the thermal dissolution of hematite, the photochemical iron(I1) formation may occur as an autocatalytic process. Near-UV light (A < 400 nm) is needed for the photochemical reductive dissolution of hematite in the presence of oxalate. A ligand-to-metalcharge-transfer transition is likely to be involved in this surface photoredox reaction, either of the iron(II1) oxalato surface complex or of the bulk hematite. More conclusive information is needed on the role of a surface complex as chromophore. For such studies, a suitable ligand would be one that forme an innersphere surface complex with surface iron(III), exhibiting a ligand-to-metalcharge-transfer band in a spectralwindow of the solid phase and being readily photolyzed. Oxygen inhibits the photochemical reductive dissolution of hematite, and the photochemical oxalate oxidation occurs at a higher rate in an aerated suspension as compared to a deaerated Suspension. Hence, hematite acta as a photocatalyst for the photochemicaloxidation of oxalate by oxygen. These experimental observations provide circumstantial evidence that in this system detachment of the reduced surface iron from the crystal lattice is the rate-determining step of the overall dissolution process. For several reasons the photochemicaliron(I1) formation may be an important process in surface waters: (i) The oxidative hydrolysis of Fe(I1) results in formation of iron(III)(hydr)oxidesthat are readily thermally dissolvedand, hence, in formation of bioavailable iron6 (ii) Fe(I1) catalyzes the thermal dissolutionof iron(II1) (hydr)oxides. (iii)The photochemical iron(I1)formation is accompanied by the photooxidation of electron donors, either biogenic or anthropogenic. Acknowledgment. We thank Werner Stumm, EAWAG, Janet G. Hering, EAWAG, and JamesJ. Morgan, California Institute of Technology, for stimulating discussions and helpful advice and Rudolf Giovanoli, University of Bern, for the characterization of the hematite phase by X-ray diffraction and electron microscopy.
