Environ. Sci. Technol. 1988, 22, 778-785
Demas, C. R.; Pereira, W. E.; Barber, L.; Updegraff, D.; Rostad, C. E.; Keck, R. J.; Chiou, C. T. Water-Resour. Invest. (U.S. Geol. Surv.), (in press). US. Environmental Protection Agency, Survey of Industrial Processing Data Task I-Hexachlorobenzene and Hexachlorobutadiene Pollution from Chlorocarbon Processes; US. EPA, Office of Toxic Substances: Washington D.C., 1975; EPA-560/3-75-003.
Moore, J. W.; Ramamoorthy, S. In Organic Chemicals i n Natural Waters, Applied Monitoring and Impact Assessment; Springer-Verlag: New York, 1984; pp 43-66. Hoese, H. D.; Moore, R. H. In Fishes of the Gulf of Mexico, Texas, Louisiana, and Adjacent Waters; Texas A&M
University Press: College Station, TX, 1977; 327 p. Grob, K. J. Chromatogr. 1973,84, 255-273. Grob, K.; Kurcher, F. J. Chromatogr. 1976,117,285-294. Barber, L. B.; Thurman, E. M.; Schroder, M. P. Open-File Rep.-U.S.
Geol. Surv. 1984, No. 84-475, 89-111.
(24) Coleman, W. E.; Munch, J. W.; Slater, R. W.; Melton, R. G.; Kopfler, F. C. Enuiron. Sci. Technol. 1983,17,571-576. (25) Chiou, C. T.; Schmedding, D. W.; Manes, M. Environ. Sci. Technol. 1982, 16, 4-10. (26) Boehm, P. D.; Quinn, J. G. Geochim. Cosmochim. Acta 1973, 37, 2459-2477. (27) Chiou, C. T.; Malcolm, R. L.; Brinton, T. I.; Kile, D. E. Enuiron. Sci. Technol. 1986, 20, 502-508. (28) Eaganhouse, R. P.; Calder, J. A. Geochim. Cosmochim. Acta 1976,40, 555-561. (29) Connor, M. S. Environ. Sci. Technol. 1984, 18, 31-35. (30) Konemann, H.; van Leeuwen, K. Chemosphere 1980, 9, 3-19.
Received for review March 23,1987. Accepted December 22,1987. T h e use of trade names in this paper is for identification purposes only and does not constitute endorsement by the U.S. Geological Survey.
Photoassisted Dissolution of a Colloidal Manganese Oxide in the Presence of Fulvic Acid T. Davld Walte," Ian C. Wrlgley,+ and Ron Szymczak Australian Nuclear Science and Technology Organisation, Lucas Heights Research Laboratories, Private Mail Bag 1, Menai, N.S.W. 2234, Australia ~
~
~~
The dissolution of a synthesized manganese dioxide by a well-characterized fulvic acid has been investigated over a range of reactant concentrations, pH, and illumination conditions, as have components of the overall dissolution process including fulvic acid and manganous ion adsorption to the oxide surface. The dissolutiotl process is satisfactotily described by an initial rapid formation of a surface-located precursor complex followed by a slower intramolecular electron-transfer step resulting in Mn(I1) formation at the oxide surface. Illumination by 365-nm light enhances the rate of electron transfer significantly, with an increase in first-order reduction rate constant from 0.31 min-l (dark) to 0.55 m i d (light) at pH 4.00 and from 0.53 min-' (dark) to 1.23 min-l (light) at pH 7.10. Depending on the affinity of the oxide surface for manganous ion, a portion of the Mn(I1) produced at the oxide surface will be rapidly released to solution, resulting in dissolution of the oxide.
Introduction In oxygenated natural waters, Mn(I1) is thermodynamically unstable and tyically forms solid manganese(II1) or manganese(1V) oxides. For example, the oxidation of manganous ion to the commonly occurring phases manganite, y-MnOOH, and birnessite, y-Mn02, i.e. Mn2++ 1/402(g)+ 3/2H20+ y-MnOOH(s)
+ 2Hf
(1)
Mn2+ + l/zOz(g)+ H 2 0 * y-MnOz(s) + 2H+ (2) proceeds with Gibbs free energy changes of -4.5 and -3.8 atm, pH 7.0, kcal (mol eT1, respectively, for P(0,) = [Mn2+]= lo4 M, and T = 25 "C [using standard electron activity (pe") values of 25.3 and 21.5 for the Mn2+/yMnOOH and Mn2+/y-Mn02half-reactions, respectively (I)]. The rate of oxide formation and type of phase produced are dependent on reaction conditions. Thus, the 'Present address: School of Chemistry, University of Melbourne, Parkville, Victoria 3052, Australia. 778
Environ. Sci. Technol., Vol. 22, No. 7, 1988
O2oxidation of Mn(1I) from oversaturated solution leads, in most instances, to the slow formation of y-MnOOH (2, 31, whereas significantly higher rates of oxide formation can be attained through surface or bacterial catalysis resulting typically in the formation of relatively disordered manganese(1V) oxides ( 4 , 5). Despite the tendency for solid manganese(II1) and manganese(1V) oxides to form in oxygenated systems, significant steady-state concentrations of solhble Mn(I1) may be generated as a result of the localized action of reductants at the oxide surface (6, 7). Natural organic substances such as the humic and fulvic acids have been implicated in this oxide dissolution process in both freshwater and seawater (6, 7) and certainly have the ability to reduce manganese(II1) or manganese(1V) oxides to the soluble manganous state. Redox potential measurements on a variety of humic and fulvic acids indicate standard electron activities (pe") in the region of 8.4 and consumption of two protons per electron transferred (B), Le., for a fulvic acid (FA) the half-reaction may be schematically written as FA(ox)
+ 2H+ + e- + FA(red)
(3) As found for the reductive dissolution of iron oxides in the presence of a fulvic acid (9),the rate of dissolution of manganese oxides is significantly enhanced on illumination (6, 7). In the case of iron oxides, the light enhancement has been associated with excitation of ligand to metal charge-transfer or internal ligand transitions within surface-located Fe(II1) complexes ( 1 0 , l l )and similar mechanisms may well apply in the manganese oxide-fulvic acid system. A semiconductor mechanism in which valence band electrons of the metal oxide are photoexcited to the conduction band creating an e- - h+ pair capable of participating in redox reactions at the particle surface has also been proposed for iron oxides (12, 13) and may be viable for some manganese oxides (14, 15). In addition, the possible role of hydrogen peroxide (produced by the photoionization of natural organic matter) as reductant has been proposed (16),i.e.
0013-936X/88/0922-0778$01.50/0
0 1988 American Chemical Society
+
FA 02- 02-+ FA+ MnO2 + HZ02 + 2H+ Mn2+ + 02 + 2H20
FA
-.
(4) (5)
The dissolution of a synthesized manganese dioxide in the presence of a well-characterized fulvic acid was investigated under a variety of conditions of reactant concentration, pH, and illumination. The dissolution results, combined with results of ancillary studies [particularly on the association of fulvic acid and Mn(I1) with the oxide surface], provide considerable insight into %e major factors controlling the dissolution process. Experimental Section Chemicals used in thii study were of analytical grade unless otherwise stated. Suwanee River fulvic acid, the standard freshwater fulvic acid of the International Humic Substances Society, was obtained from the US.Geological Survey. The extraction procedure and chemical characteristics of this material have been described elsewhere (17). AU solutions were prepared with distilled-deionized water. Preparation and Properties of Synthetic Manganese Oxide. An approximately 5-mM susaension of manganese oxide w& prepared by the dropwise addition of MnCI2 to NaMnO,, i.e.
-
(6)
gun 1. btuw image of synme~izecmanganese dbxMe showing layer sbuctue of short-term wder.
under conditions identical with thw wed by Murray (18). The prepared oxide was exhaustively washed with distilled-deionized water and the precise manganese content of the final suspension determined by atomic absorption spectrophotometry. A radiolabeled oxide was prepared in an identical fashion by spiking the MnCI2stock solution with % n C I 2 (New England Nuclear), a y-emitter with an energy of 835 keV and a half-life of 300 days. The stoichiometry of the 'cold" manganese oxide was determined by manganese reduction with oxalate and titration of excesa oxalate with permanganate (19,zO). The hydrated water content of the prepared oxide was also determined yielding a stoichiometry of MnO1.,~1.7H2O. Transmission electron micrographs of the prepared material, obtained from a JEOL l00CX instrument, indicated a single, disordered phase of small, folded leaflets. Distinct layer structure is evident in lattice images of this material (Figure 1)but is of short-term order only. The layer spacing of 7-10 A is typical of layer-structured manganese dioxides (21). Electron diffraction patterns yielded d-spacings of 2.48, 2.17, and 1.51 A, and X-ray powder spectrum, obtained on a Siemens D500 diffractometer with Co Ka source, exhibited broad peaks at 7.40, 2.43, and 1.40 A. A Fourier transform infrared spectrum of the manganese oxide, obtained from a Nicolet 10-MX spectrometer, exhibited three major peaks at 529, 1630-1630, and 3370 cm-'. The peaks above 1400 cm-I are due to hydrous components of the oxide, Le., hydroxyl and water molecules, hut the peak at 529 em-' may be attributable to Mn-0 vibrations, the position of which is characteristic of the particular manganese oxide (22,23). Three separate Brunauer-Emmett-Teller (BET)surface area analyses were performed on the synthesized cold manganese oxide. A value of 197 m2 g-' was obtained by a Carlo Erba Model 804 Sorbomatic instrument on a sample that had been evacuated at 0.1 Torr at room temperature overnight. Analysis on the same instrument after evacuation at 200 OC overnight gave a value of 186 m2g-I. A multipoint BET surface area analysis performed on a Quantachrome Quantasorb instrument yielded a value of
188 m2e;' after evacuation of the sample overnight at room temperature. The surface charge characteristies of the prepared oxide have been investigated in some detail by Murray (18). The zero p i n t of charge pH (pH,) was found to be 2.25 with electrophoretic mobilities and Na+ and K+ absorption. Dissolution Studies. All manganese oxide dissolution studies were performed in a 250-mL water-jacketed, c a p ped glass vessel, the temperature of which was maintained at 25.0 i 0.1 "C by using a Tecam C400 circulator and Tecam loo0 heat exchanger. The vessel contained a quartz plate window to enable photolysis of the suspension. All reactions were carried out in 0.05 M NaCl, and the suspensions were vigorously stirred with a magnetic stirring bar. (Preliminary studies indicated that stirring rate was not a critical determinant of dissolution rate.) Studies were performed at pH 4.00 f 0.05 and pH 7.10 0.05 with the pH maintained by continuously bubbling at 1.05 0.02% C02/air gas mixture through suspensions with bicarbonate concentrations of lodmand lo-'" M for pH 4.00 and 7.10, respectively. In experiments requiring deoxygenated conditions, a 1.05 0.02% C02/N2gas mixture was used. Studies were conducted under both illuminated and nonilluminated conditions. When required, illumination was provided by an Osram XBO 100-W mercury lamp mounted in a Rofin arc lamp and power supply system. The 365-nm Hg line was isolated by passing the beam through a narrow band-pass filter (Corning Model 5970, A, = 365 nm). The intensity of the light entering the reaction vessel was determined by potassium ferrioxalate actinometry (24) to be 96 NEcm-' mi&. Light intensity dependency studies were performed on neutral density filters (Corion Corp. QND 0.3-2 and QND 0.15-2). For the dissolution studies, an aliquot of the radiolabeled manganese oxide stock suspension was added to a dark, thermally equilibrated solution of desired pH and fulvic acid concentration. In studies requiring illumination, a shutter between the reaction vessel and the prewanned Hg lamp was opened within 5 s of colloid addition. Aliquots of the reaction mixture were removed a t predetermined intervals for analysis of total and filterable manganese.
3Mn2++ 2Mn0,-
+ 2H20
5Mno2 + 4H+
*
*
Envton. Sci. Technol.. Vol. 22. No. 7, 1988 779
Acidified H20zwas added to a 1-mL aliquot for total Mn determination, and a 2-mL aliquot was filtered through a 0.05-rm Nuclepore polycarbonate membrane for determination of filterable manganese. In some cases, manganese remaining on the filter was also determined, after digestion with acidified H202. The total, filterable and nonfilterable, manganese concentrations were determined by measuring the y-activity of the aliquots with either an LKB1280 Ultragamma Counter or a Packard 5780 AutoGamma Counter. In all cases, calibration solutions and background were counted with each set of samples. Studies of up to 4-h duration were initially conducted but shorter term studies (less than 1 h) were typically undertaken. The role of photoproduced hydrogen peroxide in the dissolution of manganese oxide under the above conditions was investigated both by the performance of studies under deoxygenated conditions and by the addition of catalase, an effective peroxide scavenger, to the reaction mixture. Adsorption of Fulvic Acid to Manganese Oxide. Manganese oxide was added to a thermally equilibrated fulvic acid solution (25.0 f 0.1 "C)and' mixed for 2 min. (The results of preliminary studies indicated that the adsorbed fulvic acid reached an equilibrium or steady-state concentration within 2 min and began to exhibit significant variability, presumably due to ongoing redox activity after about 5 min.) Multiple aliquots of the suspension were then removed and filtered. The absorbance of the filtrate was measured at 350 nm on a Hitachi U3200 spectrophotometer with 4-cm path-length cells and compared to the fulvic acid absorbance before the addition of manganese oxide. The concentration of fulvic acid adsorbed was estimated by difference. Absorbance studies were conducted over a range of manganese and fulvic acid concentrations at pH 4.00 and 7.10 under both dark and illuminated conditions. Adsorption of Mn(II)aqto Manganese Oxide. Radiolabeled manganous chloride (54MnC12)was added to a suspension of manganese oxide in fulvic acid solution, which had been mixing for 2 min. After 2 min, 2-mL aliquots were removed and filtered, and the concentration of manganous ion in the filtrate was determined by ycounting. (Preliminary studies indicated that manganous ion adsorption equilibrium was reached under these conditions in Mn'"A] = Km"d[>MnT][xHA]/(l K & [ x H A ] ) where Km = mg-' L and [ > M ~ T ]= 2.5 mg L-'.
+
7.10; for example, concentration of fulvic acid taken up from a 20 mg L-l fulvic acid solution is only 15% of that at pH 4.00. Severe difficulties in obtaining reproducible results were experienced at this pH because of the small absorbance changes associated with fulvic acid adsorption. While the precise nature of the interaction of fulvic acid with manganese oxide has not been elucidated, the interaction can usefully be modeled as a complexation reaction between manganese oxide surface sites and organic ligand metal-binding sites. Given the difficulty in describing the acid-base chemistry of both the manganese oxyhydroxide surface groups and fulvic acid [the pK's of the manganese oxide surface groups are difficult to estimate because of the highly acidic nature of the surface (18) and fulvic acids exhibit a broad distribution of acidity constants (27)], the concentrations of manganese oxide surface sites and fulvic acid metal-binding sites will be represented by [C>MnIV0H]and [CHA], respectively, where C>MnIVOH = >MnIVOH2++ >MnIVOH >MnIVO(7) and, considering the fulvic acid binding sites to be monoprotic
+
CHA = HzA+
+ HA + A-
(8)
With these definitions, a complexation reaction describing the interaction of one surface site with one fulvic acid binding site may be written as
+
C>MnIVOH + CHA + C>MnIVA H,O
(9)
where C>MnIVA = >MnIVHA++ >Mn'"A + >Mn'"OHA- + ... (10) The mass law expression corresponding to the 1:l complexation reaction shown above will be [C>MnIVA]= Kcond[C>MnlVOH][CHA] (11) where Kcondis the conditional formation constant for surface complex C>MnIVA and incorporates proton inEnviron. Sci. Technol., Vol. 22, No. 7, 1988
781
teractions with ligand and surface sites. It should of course be recognized that eq 9 represents the simplest possible model of the manganese oxide-fulvic acid interaction. The assumptions that all fulvic acid binding sites are identical and that one surface site interacts with one fulvic acid binding site are clearly oversimplications as is the neglect of the surface-active properties of fulvic acid, but in view of the complexity of the interaction and the difficulty in describing the real situation rigorously, the simple model adopted here appears reasonable. The total manganese oxide surface site concentration, [>Mn,], may be expressed as [>Mn,] = [CMnrVOH]+ [C>MnIVA]
(12)
which on substitution in the mass law expression (eq 11) yields [CHAI (13) 1 Kcond[CHA]
[C>MnIVA] = KCond[>MnT]
+
As is shown in Figure 6, a rectangular hyperbola of this mg-l form with conditional stability constant Kwnd= L and total surface site concentration [ > M ~ T(expressed ] in terms of total adsorbable fulvic acid) of 2.5 mg L-l provides a reasonable description of the adsorption of fulvic acid to 100 pM manganese oxide at pH 4.00. Note that the total surface site concentration can be expressed in molar units by recognizing that 1mg of Suwanee River fulvic acid possesses approximately 6 pmol of sites capable of binding metal ions (17). Thus, a 100 pM suspension of MnO, will exhibit a total binding site concentration of -15 pM. (This corresponds to a fulvic acid binding site concentration of r 4 . 1 sites per nm2 of oxide surface.) The adsorption data at pH 7.10 is relatively linear with little sign of saturation. In this case, a conditional stability of mg-l L is obtained by assuming constant Pond that [C>MnrVOH]N [>MnT] and utilizing the value of total surface site concentration estimated at pH 4.00. Under conditions in which the concentration of adsorbed fulvic acid [C>MnNA] is only a small fraction of the total added fulvic acid concentration [HA,], eq 13 may be approximated by an expression of the form (14) 1 + KCond[HAT]
[C>MnIVAJr KCond[>MnTJ
Such an approximation will be of particular use in estimating adsorbed fulvic acid concentrations at the low concentrations of Mn02 used in the dissolution studies. Mn(II)aqAdsorption to Manganese Oxide. The results of studies of Mn(II), adsorption to 5 pM Mn02 in the absence and presence 2 mg L-' fulvic acid at pH 4.00 and 7.10 (Figure 7) indicate continuing uptake of Mn(1Uaq, even at high sorbate concentrations, with no apparent saturation of the adsorbing surface. As indicated in Figure 7, this adsorption data can be satisfactorily described by exponential Freundlich expressions of the form
02
tMn(II)adsl = A tMn(II)aqla
(15)
As expected for adsorption to an amphoteric metal oxyhydroxide surface, the extent of cation uptake is higher at pH 7.10 than pH 4.00. The presence of 2 mg L-l fulvic acid has little effect at pH 7.10, but at pH 4.00 it leads to a small decrease in the proportion of manganous ion taken up by the oxide at each Mn(II)aqconcentration. That fulvic acid has a more marked effect on Mn(Waqadsorption at pH 4.00 than at pH 7.10 is more clearly demonstrated in Figure 8 where [Mn(II),ds] is plotted as a function of bulk fulvic acid 782
Environ. Sci. Technol., Vol. 22, No. 7, 1988
0
2
L
6
8
IMn 111I oq I.KM
Figure 7. Concentration of manganous ion take up at pH 4.00 and 7.10 by 5 pM MnO, as a function of free Mn(II)aqconcentration. Results both in the absence (solid symbols) and presence (open symbols) of 2 mg L-' fuivic acid are shown. The solid lines represent least-squares best fit exponential expressions of the form [Mn(I I)ads] = A [Mn(II), la to the organic-free data where (A, a ) = (-10.62, 0.32) and (-9.14, 8.37) at pH 4.00 and 7.10, respectively.
-
l o t
"'-----I pH 7 1 0
06
r
[FULVIC ACID
I , mg t-'
Flgure 8. Effect of fulvic acid concentration on concentration of Mn(II)aqremoved at pH 4.00 and 7.10 from a 5 pM MnO, suspension containing 2 pM added Mn(I1).
concentration. For example, at pH 7.10, 10 mg L-l fulvic acid displaces approximately 0.07 pM of the 0.72 pM Mn(1I) adsorbed in the absence of fulvic acid, but at pH 4.00 the same concentration of fulvic acid displaces 0.16 pM of the 0.32 pM Mn(I1) taken up in organic-free suspension. This increased inhibition of Mn(II)aqadsorption to the oxide by fulvic acid at the lower pH is to be expected in light of the results of fulvic acid adsorption studies reported above. The effect of fulvic acid competition for available surface sites on Mn(II),, uptake can be approximately accounted for according to the modified Freundlich expression [Mn(I1),ds] = A* [Mn(II)aql (16) where
It should be noted that solution-phase complexation of Mn(II)aqby fulvic acid could also yield the observed inhibition of Mn(II)aq adsorption. However, differential pulse polarographic studies indicate no measurable shift
in Mn(1Ijaqreduction potential or decrease in reduction peak height on addition of fulvic acid suggesting that all Mn(II)aqis present as the uncomplexed manganous ion Mn2+(aq)(28). The interaction of manganous ion with the solid manganese oxide can be represented as a complexation reaction between Mn2+(aq)and oxyhydroxide surface sites; for example, in the case of the monoprotonated surface species the reaction may be written as >MnIVOH
+ Mn2+(aq)+ H,O +
+
>Mn1VOMn110H2+ H+ (18) where >Mnw0Mn1IOH,+ represents a manganous surface species underlain by the solid manganese(1V) oxyhydroxide. The adsorption isotherm corresponding to this equilibrium expression would exhibit saturation at high sorbate concentrations (i.e., Langmuir behavior) rather than the Freundlich behavior observed, so additional processes must be invoked to account for the continuing uptake of Mn2+(aq). Formation of fresh manganese(II1) or manganese(1V) oxyhydroxide surface layers [with their own particular Mn(I1) sorbing capacity], as a result of heterogeneously catalyzed autoxidation of adsorbed manganous ion (29),most likely accounts for this continuing uptake of Mn(I1) at high sorbate concentrations. Although Freundlich expressions describe empirically the observed adsorption data, a model incorporating surface complexation of Mn(I1) with eventual precipitation of the sorbing cation (or oxidized form of the cation), together with competition for surface sites by fulvic acid, would be more satisfying. Such models have been invoked to describe Freundlich behavior in some systems (30),but in this case more definitive values of essential model parameters (particularly oxide surface and fulvic acid acidity constants) than are available are required in order to develop an unambiguous description of the adsorption process.
Discussion Reductive dissolution of metal oxides by organic reductants involves the diffusion of reactants to and products from the oxide surface. As schematized below, there are also a number of intermediate reactions at the particle surface, including adsorption of reductant resulting in precursor complex formation, electron transfer within the precursor complex resulting in formation of successor complex, breakdown of successor complex through ligand substitution, and product release (31-33). In the case of the monoprotonated surface and fulvic acid binding sites these reactions may be written as follows: adsorption >Mn'"OH
+ HA
electron transfer
Kmnd
>MnIVA
-
+ H20
(19)
ked
>MnIVA ligand exchange >MnlVOMnrlA,,
+ H20
>MnlVOMnllA,,
(20)
fast
MnrVOMn110H2++ Ao; (21)
detachment >Mn1VOMn110H2+ H+ + >MnIVOH + Mn2+(aq)+ H20 (22)
+
As indicated by the lack of effect of stirring rate on rate of dissolution, transport-controlled diffusion processes are not rate limiting under the dilute suspension conditions
used here. Attention is thus focused on the surface reactions. Photolysis is unlikely to affect the ligand exchange and detachment steps but may alter the tendency of fulvic acid to associate with the manganese oxide surface through the generation of positively charged (ionized) species, which are more readily attracted to the negatively charged oxide surface. Light may also assist the reduction of Mn(IV) and concomitant oxidation of adsorbed fulvic acid. Photoassistance of this electron-transfer process may occur by a variety of mechanisms, including (i) absorption by ligand to metal charge-transfer bands arising from the formation of inner-sphere complexes between certain fulvic acid functional groups and Mn(1V) centers at the oxide surface (34,17),(ii) excitation of n T* or T T* transitions of the fulvic acid molecule (both in solution and adsorbed to the oxide surface) with a resultant increase in electrondonating ability of the molecule (11),and (iii) generation of charge carriers (e-, h+) at the oxide surface by excitation of valence band electrons within the bulk oxide to the conduction band (14, 15). While there is little evidence that a semiconducting mechanism is appropriate for this oxide (particularly in view of the disorder apparent from electron micrographs of this manganese oxide phase), none of these possible mechanisms can be ignored without further experimental evidence. The apparent lack of significance of the proposed hydrogen peroxide mechanism is surprising in view of reports of the ability of this agent to reductively dissolve manganese oxides (16, 17, 35). Presumably, complete surface coverage of the manganese oxide by adsorbed fulvic acid prevents interaction of this Mn(1V) reductant with the oxide surface under the conditions used in this study. The reducing ability of fulvic acid under various reaction conditions can be estimated from the total concentration of Mn(I1) produced per unit time. This total concentration of Mn(I1) includes both free and adsorbed Mn(I1) species, i.e. (23) [Mn(IU+otl= [Mn(Waq1+ [Mn(II),d81
- -
[Mn(II)t,t] has been estimated for all experimental data with measured [Mn(II),] and [Mr1(11)~*]calculated by the modified Freundlich expression in eq 16. Estimates of [C>MnrVA]required in eq 16 have been obtained from eq 14 with conditional stability constants derived from the fulvic acid adsorption studies and total surface site concentrations [>MnT] = ([MnT]/lOO pM) X 2.5 mg L-l, where [Mn,] is the concentration of oxide present. These estimates of adsorbed fulvic acid will only be approximate because of the errors inherent in extrapolation of the adsorption results from high to low solids concentrations. Total Mn(I1) production rates obtained from the initial slope of [Mn(II),,] versus time data are plotted in Figures 9 and 10 as functions of total manganese oxide concentration and total added fulvic acid concentration, respectively. As discussed in the case of Mn(I1) adsorption to the oxide, the Mn(I1) retained at the oxide surface most likely undergoes relatively rapid, heterogeneously catalyzed autoxidation with resultant formation of fresh manganese(II1) or manganese(1V) oxyhydroxide (29). The contribution of the adsorbed component Mn(II)ad8 to Mn(II), is clearly very significant-particularly at pH 7.10 and high solids concentration. For example, in an illuminated 10.5 pM MnO, suspension containing 10 mg L-' fulvic acid, soluble Mn(I1) is produced at a rate of -0.7 nM s-l compared to a total Mn(I1) production rate of -7 nM s-I. Indeed, while the rate of Mn(II)aqproduction plateaus markedly on increasing manganese oxide concentration (Figure 5), the rate of total Mn(I1) production Environ. Sci. Technol., Vol. 22, No. 7, 1988 783
Table I. Values of First-Order Reduction Rate Constant k r e d and Conditional Stability Constant Kcbud Obtained from Linear Least-Squares Fits of Rectangular Hyperbolic Expressions of the Form d[Mn(II),,,]/dt = k redK o o n d[>MnT1[HATl/(l 4- Koond[HA~]) to d[Mn(II),o,l/dt versus [HAT] Data Obtained under Dark and Illuminated Conditions at pH 4.00 and 7.10
light dark
[MANGANESEI,pM
Figure 9. Effect of manganese oxide concentration on estimated rate of reduction of MnO, suspended in 10 mg L-' fuivic acid at pH 4.00 and 7.10. c
x
pH 4.00 kred,min-' KOond,mg-' L 0.55 10-067 0.31 10-0l9
pH 7.10 kred, min-' Kcand, mg-' L 1.23 10-1 44 10-164 0.53
Under these conditions, the adsorption process may be considered to be a preliminary equilibrium step preceding electron transfer. Assuming that the concentration of the precursor complex reaches a steady state, eq 25 may be written 37
:: w c (L 4 2
0 c
1
s z (L
5
0
bpH710
0
20
10 15 IFULVIC A C I D ] , mg L - '
19
'
' 4I
I
20 30 [FULVIC A C I D ] . m y C'
40
Figure I O . Effect of fulvlc acid concentration on estimated rate of reduction of 5 pM MnO, under dark and illuminated conditions at (a) pH 4.00 and (b) pH 7.10. Solid lines represent least-squares best fit rectangular hyperbolae of form d[Mn(II),,]ldt = k,, KCond[>MnT]KCOnd[HAT]) where the fitting parameters kred and KcoM [HA,]/(l ihave values shown in Table I.
increases linearly with increase in oxide concentration (Figure 9). As found for the dependency of Mn(1Uaq production rate on fulvic acid coilcentration, the rate of Mn(II), production decreases on increasing concentration of fulvic acid. The processes leading to the production of Mn(II), may be simply represented (for the monoprotonated surface and fulvic acid binding sites) as >MnIVOH
+ HA G >MnIVA p n d
kred
Mn(II)tot
(24)
where kredis the first-order rate constant for the reduction step, i.e.
Studies of the interaction of a variety of organic reducing agents with dilute suspensions of manganese oxides indicate that ligand adsorption/desorption rates are typically rapid compared to the rate of electron transfer (7, 36). 784
Environ. Sci. Technol., Vol. 22, NO. 7, 1988
Since the total surface site concentration, [>MnT], is directly proportional to the concentration of added rnanganese oxide, the proportionality between rate of production of Mn(1Ubt and added oxide concentration observed in Figure 9 is to be expected from eq 26. Best fit rectangular hyperbolas of the form shown in eq 26 have been obtained for the relationship between reduction rate and fulvic acid concentration (Figures 9 and 10). These fits yield the estimates af first-order reduction rate constants and conditional stability constants given in Table I. Very good fits were obtained at pH 7.10, but some deviation from the rectangular hyperbolic model is observed at pH 4.00. This is possibly due to oversimplifications in the method of accounting for competition between the organic acid and manganous ions for available surface sites (a problem that is particularly severe at the lower pH). The conditional stability constants obtained from the dissolution data are significantly higher than those from the adsorption studies, where 20-fold higher concentrations of manganese oxide were used. This is not particularly surprising since there is evidence of marked increases in the adsorption density of naturally occurring organic matter on metal oxides when oxide concentration is decreased (26). The formation constants reported in Table I exhibit both the expected pH dependency (higher at lower pH) and the previously observed increase on illumination. The first-order reduction rate constants reported in Table I are ca. 2 times higher at pH 7.10 than at pH 4.00, suggesting that the fulvic acid has more reducing ability at the higher pH. Although such an effect of pH might be expected from the deduced half-reaction for fulvic acid (eq 3), the pH dependence of the MnZ+/Mn02couple must also be taken into account. In fact, the stoichiometry of both electron donor and electron acceptor processes would be expected to involve two protons per electron transferred with zero net production or consumption of protons. The fact that there is only ca. 2-fold change in the reduction rate with a 1000-fold difference in [ H + ] suggests that the two strong pH dependencies do essentially cancel out. The overall quantum yields for photoassisted reduction of MnO, to Mn(II)bt by 365-nm light at pH 4.00 and 7.10 can be assessed from the difference between the rate constants obtained under illuminated and dark conditions at each pH. Assuming that Mn(I1) is produced as a result of light absorption by the precursor complex C>MnIVA and that the optical density tL[C>MnTVA]within the
(4) Sung, W.; Morgan, J. J. Geochim. Cosmochim.Acta 1981, 45, 2377.
reaction suspension is small, we may write (34) ' d [ C>Mn'"A]
( [C>MnIVA1iight- [C>Mn'VAldar~)
dt = k!i$t - k$$k = 2.3O3c#JIO(S/v)~L
(27)
where e is the extinction coefficient of the precursor complex, L is the optical pathlength within the reaction vessel, I, is the incident light intensity, S is the front window area, Vis the volume of suspension, and c#J is the overal quantum yield. Assuming that the extinction coefficient of the precursor complex C>MnIVA is similar to that of the uncomplexed fulvic acid (a reasonable assumption provided photoassisted electron transfer arises predominantly as a result of absorption by internal ligand bands), we obtain overall quantum yields for the photoassisted production of Mn(II), of 2.0 X lo4 and 4.8 X lo4 at pH 4.00 and pH 7.10, respectively. Conclusions Fulvic acid is a mild reducing agent with the ability to induce dissolution of manganese dioxides. The dissolution process is satisfactorily described by an initial rapid formation of a surface-located Mn(1V)-fulvic acid precursor complex followed by a slower intramolecular electrontransfer step resulting in Mn(I1) formation at the oxide surface. Depending on the affinity of the oxide surface for manganous ion, a portion of the Mn(1I) so produced will be released rapidly to solution. Photoproduced hydrogen peroxide does not play a significant role in the dissolution process under the conditions used in this study. The rate and extent of oxide dissolution are very dependent on suspension acidity since both the degree of association of the reductant molecules with the oxide surface and the tendency for produced Mn(I1) to be released to solution are strongly pH dependent. Thus, low pH causes an increase in the rate of reductive dissolution by favoring both the adsorption of fulvic acid to the oxide surface and the release of Mn(I1) to solution. Illumination of manganese dioxide suspended in fulvic acid solution enhances the rate and extent of oxide dissolution significantly. This enhancement is caused by an increased tendency of fulvic acid molecules to associate with the oxide surface on illumination (possibly due to photoionization of organic molecules) and by the creation of photoexcited precursor species, which exhibits more facile electron transfer than occurs thermally. The results of studies conducted at pH 4.00 and 7.10 indicate that light is particularly effective in enhancing the rate of Mn(I1) production at the higher pH. The overall quantum yield for Mn(I1) production by 365-nm light at pH 7.0 is more than double that at pH 4.00,though the effect of this difference on rate of dissolution is masked by the greater tendency for manganous ions to be retained a t the oxide surface at the higher pH. Acknowledgments The assistance of J. D. Smith, University of Melbourne, in the early stages of this project is gratefully acknowledged. Registry No. MnOz, 1313-13-9. Literature Cited (1) Bricker, 0. Am. Mineral. 1965,50, 1296. (2) Stumm, W.; Giovanoli, R. Chimia 1976, 30, 423. (3) Hem, J. D.; Lind, C. J. Geochim. Cosmochim. Acta 1983, 47, 2037.
(5) Diem, D. Ph.D. Thesis, Swiss Federal Institute of Technology (ET"), Zurich, 1983. (6) Sunda, W. G.; Huntsman, S. A,; Harvey, G. R. Nature (London)1983, 301, 234. (7) Stone, A. T.; Morgan, J. J. Environ. Sci. Technol. 1984,18, 617. (8) Waite, T. D. In Humic Substances: Structure and In-
teractions; Hayes, M. H. B., MacCarthy, P., Malcolm, R. L., Swift, R. S., Eds.; Wiley: New York (in press). (9) Waite, T. D.; Morel, F. M. M. Environ. Sci. Technol. 1984, 18, 860. (10) Waite, T. D.; Morel, F. M. M. J. Colloid Interface Sci. 1984, 102, 121. (11) Waite, T. D.; Torikov, A,; Smith, J. D. J. Colloid Interface Sci. 1986, 112, 412. (12) Faust, B. C.; Hoffmann, M. R. Environ. Sci. Technol. 1986, 20, 943. (13) Leland, J. K.; Bard, A. J. J. Phys. Chem. 1987, 91, 5076. (14) Shuey, R. T. Dev. Econ. Geol. 1975,4, Chapter 23. (15) Harriman, A.; Richoux, M.-C.; Christensen, P. A.; Mosseri, S.; Neb, P. J. Chem. SOC., Faraday Trans. 1 1987,83,3001. (16) Sunda, W. G.; Huntsman, S. A. EOS, Trans Am. Geophys. Union 1983, 64, 1029. (17) Thurman, E. M.; Malcolm, R. L. In Aquatic and Terrestial
Humic Materials; Christman, R. F., Gjessing, E. G., Eds.; Ann Arbor Science: Ann Arbor, MI, 1983; Chapter 1, pp 1-23. (18) Murray, J. W. J . Colloid Interface Sci. 1974, 46, 357. (19) Treadwell, F. R.; Hall, W. T. In Analytical Chemistry, 9th ed.; Wiley: New York, 1942; p 559. (20) Murray, J. W.; Balistieri, L.; Paul, B. Geochim. Cosmochim. Acta 1984, 48, 1237. (21) Giovanoli, R.; Burki, P.; Giuffredi, M.; Stumm, W. Chimia 1975, 29, 110. (22) Swinkels, D. A. J.; Fredericks, P. M.; Anthony, K. E. Proc.-Electrochem. SOC.1985, No. 85-4, p 158. (23) Potter, R. M.; Rossman, G. R. Am. Miner. 1979,64, 1199. (24) Murov, S. L. Handbook of Photochemistry;Dekker: New York, 1973. (25) Gorichev, I. G.; Kipriyanov, N. A. Russ. Chem. Rev. (Engl. Trans.) 1984, 53, 1039. (26) Davis, J. A. Geochim. Cosmochim. Acta 1982, 46, 2381. (27) Dzombak, D. A.; Fish, W.; Morel, F. M. M. Environ. Sci. Technol. 1986, 20, 669. (28) Batley, G. E., Lucm Heights Research Laboratories, Sydney, personal communication, 1987. (29) Davies, S. H. R. In Geochemical Processes at Mineral
Surfaces;Davis, J. A., Hayes, K. F., Eds.; ACS Symposium Series 323, American Chemical Society: Washington, DC, 1986; Chapter 23. (30) Farley, K. J.; Dzombak, D. A.; Morel, F. M. M. J. Colloid Interface Sci. 1985, 106, 226. (31) Stone, A. T. In GeochemicalProcesses at Mineral Surfaces; Davis, J. A,, Hayes, K. F., Eds.; ACS Symposium Series 323, American Chemical Society: Washington, DC, 1986; Chapter 21. (32) Waite, T. D.; Torikov, A. J. Colloid Interface Sci. 1987,119, 228. (33) Stone, A. T.; Morgan, J. J. In Aquatic Surface Chemistry:
Chemical Processes at the Particle- Water Interface; Stumm, W., Ed.; Wiley: New York, 1987; Chapter 9. (34) Balzani, V.; Carassiti, V. In Photochemistry of Coordination Compounds; Academic: London, 1970. (35) Bard, S.; Lume-Pereira, C.; Janata, E.; Henglein, A. J. Phys. Chem. 1985,89,5779. (36) Stone, A. T.; Morgan, J. J. Environ. Sci. Technol. 1984, 18, 450. (37) Espenson, J. H. In Chemical Kinetics and Reaction Mechanisms; McGraw-Hill: New York, 1981; Chapter 4.
Received for review December 29, 1986. Revised manuscript received November 20,1987. Accepted January 22,1988, This work was supported, in part, through the Australian Marine Sciences and Technologies Grants Scheme under Grant 8510933. Environ. Sci. Technol., Vol. 22,
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