Langmuir 1992,8, 95-103
95
Reductive Dissolution of Manganese(II1,IV) (Hydr)oxides by Oxalate: The Effect of pH and Light Aglaia G. Xyla,? Barbara Sulzberger,**tGeorge W. Luther, III,t Janet G. Hering,? Philippe Van Cappellen,? and Werner Stummt Institute for Water Resources and Water Pollution Control ( E AW A G ) , Zurich, Swiss Federal Institute of Technology (ETH), Zurich, Switzerland, and College for Marine Studies, University of Delaware, Lewes, Delaware 19958 Received May 23, 1991. In Final Form: September 9, 1991
The dissolution of y-MnOOH and 8-MnO2 in the presence of oxalate was studied at various pH values, in the dark and under the influence of light. Unlike iron(II1) (hydr)oxides,y-MnOOH and @-MnOzare reductively dissolved on a time scale of hours in the dark at low pH. A strong pH dependence of the rate of reductive dissolution of both manganese (hydr)oxides is observed; the rate of reductive dissolution decreases with increasing pH. This behavior may be interpreted, in part, in terms of the pH dependence of oxalate adsorption on the surface of the manganese(II1,IV)(hydr)oxides. A plateau in the dissolution rate of y-MnOOH is observed at high (total) oxalate concentrations, consistent with direct dependence of the rate of the reductive dissolution of the manganese(II1,IV) (hydrloxides on the concentration of surface-boundoxalate. At any given pH, the reductive dissolution rate per unit surface area is higher for y-MnOOH than for O-MnOz. Light has only a weak catalytic effect on the dissolution kinetics of the two manganese (hydr)oxidesstudied, in contrast to iron(II1) (hydr)oxides.
Introduction Manganese is a widely distributed element of considerable importance in natural aquatic systems. It is an essential plant nutrient and many manganese (hydr)oxides exhibit a significant adsorption capacity for numerous trace elements. In natural waters manganese is found in the +I11 and +IV oxidation states as solid manganese(II1,IV) (hydr)oxides and in the +I1 oxidation state as Mn(1Uaq, Mn(OH)2, and MnS and in association with many Under the conditions usually encountered in natural waters, the reduction of Mn(II1,IV)is accompanied by dissolution and the oxidation of Mn(I1)by precipitation. As a consequence of reductive dissolution of manganese(II1,IV) (hydr)oxides, adsorbed compounds are released into the water and, as a consequence of oxidative hydrolysis of Mn(II), reactive compounds may be adsorbed and/ or coprecipitated a t the freshly formed manganese(II1,IV) (hydr)oxide surface. Hence, the redox cycling of manganese is coupled to the geochemical cycling of other In lake waters, 7-MnOOH has been found to be the final product of the sequence of oxidation of Mn2+ by oxygen with MnaOI as an intermediate.' At the sediment-water interface, manganese(1V) oxides may be formed by oxidation of upwardly diffusing Mn2+.lv516 In oceanic surface waters, enrichments in dissolved manganese concentration are widely o b ~ e r v e d . ~Such -~ t Swiss Federal Institute of Technology.
University of Delaware. (1)Stumm, W.; Morgan, J. J. Aquatic Chemistry: An Introduction Emphasizing Chemical Equilibria in Natural Waters; Wiley-Interacience: New York, 1981. (2)Crerar,D.A.;Cormic,R. K.;Barnes,H.L. InGeologyandGeochemistry of Manganese; Varentsov, I. M., Grasselly, G. Y., Eds.; Akadbmiai Kiado: Budapest, 1980. (3)Burns, R.G.;Burns, V. M. InMarineManganeseDeposits; Glasby, G. P., Ed.; Elsevier Oceanography Series; Elsevier Science Publishers Co.: Amsterdam, 1977. (4)Murray, J. W. Geochim. Cosmochim. Acta 1975,39,505. (5)Giovanoli, R.; Biirki, P.; Giuffredi, M.; Stumm, W. Chimia 1975, 29,517. (6)Stumm, W.; Giovanoli, R. Chimia 1976,30,423. (7)Bender, M. L.; Klinkhammer, G. P., Spencer, D. W. Deep-sea Res. 1977,24,799. (8)Landing, W. M.; Bruland, K. W. Earth Planet. Sci. Lett. 1980,49, 45.
enrichments are usually attributed to in situ photochemical reductive dissolution of manganese(II1,IV) (hydrloxides or input from external sources. External sources of manganese include riverine and atmospheric inputs,&'O upwelled Mn-rich deep ~ e a w a t e r and , ~ manganous ion dispersion from reducing sediments into the overlying water c0lumn.~9~ Additionally, the persistence of relatively high concentrations of dissolved Mn(I1)found in the photic zone of surface waters may be due in part to the slow kinetics of Mn2+oxidation, particularly below pH = 8.5, in the absence of catalysis by mineral surfaces or micro~rganisms.l~-'~ Several studies have investigated the effect of organic reductants on the dissolution of manganese(II1,IV) (hydr)oxides.16-20These studies have shown that manganese(II1,IV) (hydr)oxides are reductively dissolved by a variety of aromatic and nonaromatic compounds in the absence of light. Photodissolution studies using natural fulvic acid in both freshwater and seawater21p22 have shown that light enhances the dissolution of manganese(II1,IV) (hydr)oxides. In order to clarify the roles played by light and organic ligands, we have investigated the reductive dissolution of (9)Statham, P. J.; Burton, J. D. Earth Planet. Sci. Lett. 1986,79,55. (10)Statham, P. J.; Burton, J. D. Geochim. Cosmochim. Acta 1988, 52,2433. (11)Morgan, J. J.; Stumm, W. Proceedings ofthe Second Conference on Water Pollution Research; Pergamon: Elmsford, NY, 1964. (12) Diem, D.; Stumm, W. Geochim. Cosmochim. Acta 1984,48,1571. (13)Davies, S. H. R.; Morgan, J. J. J. Colloid Interface Sci. 1989,129, 63. (14)Davies, S. H. R. In Geochemical Processes at Mineral Surfaces; Davis, J. A., Hayes, K. F., Eds.; American Chemical Society: Washington, DC, 1986; Chapter 23. (15)Wehrli, B. In Aquatic Chemical Kinetics: Reaction Rates of Processes in Natural Waters; Stumm, W., Ed.; Wiley-Interscience: New York, 1990; Chapter 11. (16)Stone, A. T.;Morgan, J. J.; Enuiron. Sci. Technol. 1984,18,450. (17)Stone, A. T.;Morgan, J. J.; Enuiron. Sci. Technol. 1984,18,617. Morgan, J. J. In Aquatic Surface Chemistry;Stu", (18)Stone, A.T.; W., Ed.; Wiley-Interscience: New York, 1987;Chapter 9. (19)Stone, A. T.Geochim. Cosmochim. Acta 1987,51,919. Ulrich, H. J. J. Colloid Interface Sci. 1989,132,509. (20)Stone, A. T.; (21)Sunda, W. G.; Huntsman, S. A.; Harvey, G. R. Nature 1983,301,
-- -. 92A
(22)Waite, T.D.; Wrigley, I. C.; Szymczak, R. Enuiron. Sci. Technol. 1987,18, 860.
0 1992 American Chemical Society
96 Langmuir, Vol. 8, No. 1, 19.92
Xyla et al.
both a mmganese(IT1) hydroxide, r-MnOOY, and a manganese(1V) oxide, P-MnOz, in simple model systems with oxalate as the reductant. The ligand/alectron domr, oxalate, is used as an analogue for natural organic acids.lg
Experimental Section Prepcsration and ChRrmtedmtion of the Manganese(II1,IV) (Hydr)oxides. Manganite (y-MnOOY) WPS prewred according to the method of L e ~ e n b e r g e r .Three ~ ~ hundred mL of NM3 (70% ) was added under vigorous stirring to a mixture of 20.4 mL of H202 (30%) and 11mL of 0.06 M MnS04 solution. The solution was the0 quickly heeted to boiling point and lrmt for 6 h at 95 "C. The procliact was collected (still hot) by filtration and washed repeatedly with a total of 1000 m L of hot water to remove all NH4+ and S042-.The product was dryed in vacuo a t room temperature over P2O5 for about 48 h. It was characterized by X-ray powder diffraction and by electron transmission microscopy (Figure la). The electron micrographs show needles of approximately 350 X 40 X 25 nm. The specific surface area as determined by BET measurements was 30.5 m2 g-l. The concentration of exchangeable surface sites was determined by fluoride adsorption measurements as 2.5 X 10"' mol g1.24 From the specific surface area and the total exchange capacity measurements we calculated a surface density of active sites equal to about five OH- groups per square nanometer. Sodium birnessite (N&Mn14027*9H20)was prepared following the method of Staehli.25 Two hundred milliliters of 0.5 M Mn(NO& was added to 250 mL of NaOH 5.5 M; 0 2 was then bubbled through the resulting Mn(OH)2 suspension a t a rate of 2.5 L min-l. After 4 h oxygenation was finished and the black precipitate was thoroughly washed with cold deionized water and dried in a vacuum desiccator with PZO5. The product (Figure lb) was characterized as described for y-MnOOH. The BET specific surface area was 5.2 m2 rl. The third oxide phase, B-MnO2 (Merck), was obtained commercially and used as received (Figure IC).The X-ray diffraction pattern exhibits the characteristic reflections of pyrolusite. The BET surface area was 214 m2 gl. The surface acidity constants for manganite, corresponding to the following equilibrium reactions at the oxide surface >Mn"'OH:
>Mn"'OH
+ H+
Kat
(1)
>Mn"'OH e >Mn"'O- + H+ Ka: (2) were determined by acid-base titrations of CO2-free manganite suspensions according to the method described by Stumm and Morgan1and Sigg.% The intrinsic protolysis constants were found to be pKalint= 5.72 and pKazint= 6.66. The corresponding point of zero charge is pH,, = 6.2, and double layer capacitanceobtained from the slope of the acidic branch of the titration curve is K = 1.02 f 0.1 F m-2 (0.01 M).27928 Dissolution Experiments. All solutions were prepared with deionized distilled CO2-free water and analytical grade reagents. The experiments were carried out in the absence of light in thermostated batch reactors (25 & 0.1 "C) a t a constant ionic strength of 0.05 M (NaC104). The suspensions were purged with nitrogen that had previouslypassed a Jones reductor and vigorously stirred throughout the dissolution experiments in order to prevent settling of the particles. The reaction volume was 300 mL. The oxide particles (0.33 g L-l) were pretreated for 18 h at the experimental pH value and ionic strength, prior to the addition of the ligand. The ligand was added from a solution of Na2C204 with trace concentrations (ca. 1% ) of radiolabeled (14C) oxalic acid. The pH was kept constant by a pH-stat through the addition of HClO4. The pH-combination electrode was calibrated beforehand using pH buffer solutions. Separation of dissolved and particulate manganese was achieved with 0.2 pm membrane (23)Giovanoli, R.;Leuenberger, U. Helu. Chim. Acta 1969,52,2333. (24)Faust, B. C. Ph.D. Thesis, California Institute of Technology, Pasadena, CA, 1985. (25) Giovanoli, R.; Staehli, E.; Feitknecht, W. Helu. Chim. Acta 1970, 53,453. (26)Sigg, L. Ph.D. Thesis, ETH Ziirich, No. 6417,1979. (27) Muller, B.; Sigg, L. Aquat. Sci. 1990,52,75. (28)Muller, B.;Sigg, L. J. Colloid Interface Sci., in press.
Figure 1. Transmission electron micrographs of y-MnOOH (a, top), N&Mn140~9H20(b, middle) and B-Mn02 (c, bottom) crystals. filters, and the concentration of dissolved manganese was determined by flame atomic absorption spectrophotometry. Colorimetric tests, using the o-tolidine method,29 indicated the absence of oxidation states higher than +I1 for manganese in solution. Oxalate concentrations in the filtered solutions were determined directly by scintillation counting.30 In addition, the C02 produced by the oxidation of oxalate was precipitated as BaC03 in a Ba(OH)2 solution. At the end of an experiment the amount of unreacted Ba(OH)2 was titrated with HCl of appro(29)Stumm, W.;Morgan, J. J. J. Am. Water'Works Assoc. 1965,107. (30)Banwart, S.Ph.9. Thesis, ETH Ziirich, No 8934, 1989.
Reductive Dissolution of MnOOH and MnO2 by Oxalate
Langmuir, Vol. 8, No. 1, 1992 97
priate concentration. The amount of C02 produced in the experiments was also determined independently by scintillation counting of the collected BaCQ suspension. The two methods for measuringC02 production were found to be in good agreement. Control experiments in which oxide particles were suspended in the same ionic medium in the absence of oxalate showed no measurable loss of manganese to the reaction vessel walls. For the photochemical experiments, an experimental setup was used as described by Siffert and Sul~berger.~~ The experiments were carried out with white light from a 1000-W highpressure xenon lamp (OSRAM) that was filtered by the boxtom window of the Pyrex glass vessel, which acts as a cutoff filter (A112 = 350 nm). The incident light intensity was IO zz 0.15 kW m2. Speciation Calculations. The computer speciation program MICROQL,S~ based on the constant capacitance model for adsorption on charged surfaces, was employed to calculate the speciation of oxalate at the y-MnOOH surface. The experimentally determined intrinsic surface acidity constlints for y-MnOOH served as input to the program. Because the adsorption constant of oxalate on y-MnOCiH has not yet been experimentally determined, we used the corresponding intrinsic equilibrium constant for oxalate adsorption on goethite (a-FeOOH)33 >FeOH + H C 2 0 i
>FeC,OL
+ H20
pKlint= 7.6 (3)
500
1". [oxalate]
' 0
.
'
" 100
. ' u
'
'
200
2MnOOH(s) MnO,(s)
+ CO, :-
+ CO, :-
+ 6H'
-
9 4H+
-
2Mn2++ 2C0, -k 4pI[,o (4)
+
Mn2+ 2C0,
+ 2H,O
(5)
where the speciation of oxdate correspondsto its addition from a solution of Na2C204 (see Experimental Section). At pH 4, close to the second pKa of oxalate, some of the added oxalate will be protonated to form the species HC~OQ. If, however, most of the added oxalate is consumed in the reaction, the protonation of oxalate will have little effect on the overall stoichiometry expressed above. Some release of Mn2+into solution is observed prior to the addition of oxalate. This release Q C C U P ~during the initial equilibration of the solid at the pH of the dissolution experiments (see Experimental Section). The extent of Mn2+release during pretreatment decreases with increasing pH and is greater for the manganese(Bi1) hydroxide than for the manganese(1V) oxide at corresponding pH (31) Siffert, C.; Sulzberger, B. Langmuir 1991, 7, 1627. (32) Westall, J. C. Internal report, EAWAG. (33) Balistrieri,L. S.;Murray,J. W. Geochim. Chosmochim. Acta 1987, 51, 1151. (34) Balistrieri,L.S.;Murray, J. W. Geochim. Cosmochim. Acta 1982, 46, 1041.
400
Time (min)
b
The latter choice is justified because of the similarity in surface area and density of the functional groups between goethite and y-Mn00H.33*34
Results Both manganese(111) hydroxide and mmganese(1V) oxide, y-MnOOH and p-Mn02, undergo thermal reductive dissolution in the presence of oxdate at low pH (Figure 2). Comparison of oxdate and dissolved Mn2+concentrations during the dissolution reaction demonstrates that the ratio of oxalate consumed to Mn2+produced is 1:2 for the manganese(II1) hydroxide (Figure 3a) and 1:l for the manganese(1V) oxide (Figure 3b). Protons are also consumed in the reaction of manganese(II1,IV) (hydr)oxides with oxalate; the addition of acid was required to maintain constant pH during the reaction. For the manganese(I1I) hydroxide, three protons are consumed per Mn2+released (Figure 3c). For the manganese(1V)oxide, the ratio is 4:l (Figure 3d). Thus the following stoichiometric reactions can be written
300
0
I
0
Time (min)
Figure 2. Concentrations of dissolved Mn(I1) and oxalate as a function of time upon reductive dissolution of y-MnOOH (a) and B-MnOz (b) in the presence of oxalate in the dark. Initial oxalate concentration: 1mmol L-l; pH = 4.0.
(see Table I). The concentrations of dissolved Mn2+ produced by pretreatment of the solid, however, were always much lower than the final Mn2+concentrations after reaction of the (hydr)oxides with oxalate (Figure 4). The dependence of the rate of reductive dissolution of the mmigmese(I'I1)hydroxide on the (initial)concentration of oxalate (pH 4) is illustrated by Figure 5, which shows the concentrations of Mn2+over time with (initial) oxd a t e concentrations of 10, 1,and 0.05 mM. Only a slight (approximate30 % ) increase in dissolution rate is observed with the 10-fold increase in (initial) oxalate concentration in the higher concentration range, while in the lower concentration range, the dissolution rate depends strongly on the oxalate concentration. These observations are consistent with a direct dependence of the dissolution rate on the surface concentration of oxalate. The relative insensitivity of the dissolution rate to the oxalate concentration in the higher concentration range can be attributed to saturation of the surface. The rate of y-Mn00H dissolution by oxalate exhibits ti marked dependence on pH as shown in Figure 6. Similarly, the rate of reductive dissolution of the P-Rln02 decreases rapidly with increasing pH. At any given pH, the reductive dissolution rate (normalized for the surface area of the (hydr)oxides) is higher for the manganese(II1)
Xyla e t al.
98 Langmuir, Vol. 8,No.1, 1992
\
200
4
I
400
200
0
600
3000
d
2000
El W
+
3
e 1000
Figure 3. Reductive dissolution of the solid manganese phases in the presence of oxalate in the dark at pH = 4.0with initial oxalate concentration 1mmol L-l: (a) dissolved oxalate concentration as a function of the concentration of Mn2+ produced with y-MnOOH; (b) dissolved oxalate concentration as a function of the concentration of MnZ+produced with b-Mn02; (c) concentration of consumed H+asa function of the concentration Mn2+produced with y-MnOOH; (d)concentration of consumed H+as a function of the concentration of Mn2+ produced with b-MnOz. The Mn2+ and H+ concentrations shown include the Mn2+produced or H+ consumed during preconditioning as well as during reaction with oxalate (compare Figure 4). Hence, the intercept of the x axis in parts c and d corresponds to the Mn2+concentration at time zero with regard to oxalate addition. Table I. Mn2+Concentrations after Pretreatment of 7-MnOOH and B-MnOz at Different pH Values Prior to the Addition of the Ligand [Mn2+l,pmol L-l PH y-MnOOH P-MnOz 4 150 90 5 6
a5 35
60
hydroxide t h a n for t h e manganese(1V) oxide (see Figure 8). Possible explanations for t h e observed p H effect are given in the next section. In the presence of light, a moderate acceleration over t h e rate of t h e thermal, reductive dissolution of y-MnOOH and @-Mn02by oxalate is observed (Figure 7). A t pH 4, the dissolution rate in t h e light is approximately 20% faster for t h e manganese(II1) hydroxide a n d a factor of 2 faster for t h e manganese(1V) oxide t h a n t h e dissolution in t h e dark. A somewhat larger effect of light is observed at higher pH. T h e combined effects of pH and light on t h e initial rate of reductive dissolution of y-MnOOH a n d @MnO2 are summarized in Figure 8. Also shown on t h e figure are rate data from two experiments carried out at
before addition of oxalate
0
after addition of oxalate
8
e 1200' c? E 800
i
Q
"1 0
300
6w
900
1200
1500
Time (min)
Figure 4. MnZ+formation from y-MnOOH in the dark at pH 4.0. The origin of time corresponds to the introduction of the solid in oxalate-free solution. After approximately 18 h of pretreatment, oxalate is added to the suspension. Initial oxalate concentration was 1 mmol L-l.
Langmuir, Vol. 8, No. 1, 1992 99
Reductive Dissolution of MnOOH and MnOz by Oxalate
Time (min) Figure 5. Effect of the initial dissolved oxalate concentration on the rate of reductive dissolution of y-MnOOH in the dark at pH = 4.0. 2000
isotropy measurements that both manganese(II1) hydroxide and manganese(1V) oxide also contain a small fraction of MII(II)?~This may explain the release of Mn2+observed prior to the addition of oxalate. In the case of the manganese(II1) hydroxide, the disproportionation reaction
/
1500
2Mn3++ 2H,O s MnO,(s)
--
0 0
7.0
40
60
pH=6 80
100
Time (min) Figure 6. Effect of pH on the rate of reductive dissolution of 7-MnOOH in the dark. Initial oxalate concentration was 1mmol L-1.
different pH values in the dark with NaMn1402~9HzO (birnessite). The rates of the reductive dissolution of birnessite are comparable to those of @-MnO2under similar conditions. The followingmajor observations are made in this study: (i) The reductive dissolution of the manganese(II1,IV) (hydr)oxides by oxalate exhibits saturation suggesting dependence on the surface concentration of adsorbed oxalate. (ii) The rates of the reductive dissolution of yMnOOH and @-Mn02are strongly pH dependent; the rates increase with decreasing pH. (iii) Under the same conditions the rate of reductive dissolution of y-MnOOH is much higher than that of B-MnOz. Consequently, measurable dissolution rates for y-MnOOH are obtained over a larger pH range than for P-Mn02. (iv) Light has a weak catalytic effect on the dissolution of y-MnOOH and @-MnO2.
Discussion Release of Mn2+during Oxide Pretreatment. It has
been shown on the basis of magnetic predictions and an-
+ Mn2++ 4H'
log K = 19.1 (6) may also contribute to Mn2+ release. Even at pH 4, however, the concentrations of dissolved Mn2+resulting from the oxide pretreatment amount to only a few percent (=8%) of the total manganese dissolved after reaction with oxalate. Role of Disproportionation in the Reductive Dissolution of Manganese(II1) Hydroxides by Oxalate. The reductive dissolution of y-MnOOH could in principle proceed via either of the following two pathways: (i) reduction of a Mn(II1) surface center followed by detachment of Mn(II), according to the overall stoichiometry shown in eq 4, or (ii) disproportionation of a surface Mn(111)center forming one dissolved Mn2+ion and one surface MnO2 group which then is reductively dissolved. The second mechanism can be written as 2MnOOH + 2H+ MnO,
-
+ C,042-+ 4H'
Mn2'
-
+ MnO, + 2H,O
+
Mn2+ 2C0,
+ 2H,O
(7)
(8)
From the point of view of the overall stoichiometry both pathways are possible since in each case two Mn2+ are formed and six protons are consumed per oxalate oxidized (see Figure 3a,c). The pathway involving disproportionation can, however, be excluded on the basis of two arguments. First, the rate of dissolution, normalized for oxide surface area, is faster for the manganese(II1) hydroxide than for the manganese(1V)oxide (see Results). This makes it unlikely that MnO2 formed by disproportionation is an intermediate in the overall manganese(II1) hydroxide dissolution. Rather, the kinetic data suggest that the reductive dissolution of Y-MnOOH proceeds through a one-electron transfer reaction, Mn(II1) + e-
-
(35)Krishnan, K. S.; Banerjee, S. Trans.Faraday SOC.1939,35, 385.
Xyla et al.
100 Langmuir, Vol. 8, No. 1, 1992
no overgrowth of MnO2 was observed on the surface of the manganese(II1) hydroxide. Dependence of Reductive Dissolution Rates on Surface Oxalate Concentrations. As mentionedabove, the observed dependence of dissolution rates on the (initial) oxalate concentration suggests that the dissolution rates are directly related to the surface concentration of oxalate. This is a very common observation in dissolution studies of sparingly soluble minerals in the presence of ligand^.^^-^^ The general mechanism for the reductive dissolution of a metal (hydrloxide includes the following steps:18,42s43 (1) surface complex formation between the adsorbed ligand and the surface metal center; (2) electron transfer within the surface complex, resulting in a reduced surface metal center and the oxidized ligand; and (3) detachment of the reduced surface metal center. These individual steps are given in eqs 9-11 for the case of the reductive dissolution of the manganese(II1) hydroxide by oxalate
in the light 1500
1OOo:
8E
E
500 :
40
20
0
80
Time (min)
>Mn"'OH,+ >Mn"'C,O;
+ C204,-'*
-
>Mn"'C,O,-
>Mn(II)
+ H,O
+ C0,'- + CO,
(9) (10)
+
0
10
20
40
30
50
Time (min) Figure 7. Effect of light on the rate of reductive dissolution of y-MnOOH (a) and O-MnOz (b) at pH = 4.0. Initial oxalate concentration was 1 mmol L-l. -5
4 I 4 4
A
0
e
i
A
A 3
4
5
6
>Mn(II)- > Mn(II)(aq) (11) where > denotes the surface lattice of the manganese(II1) hydroxide and >MnW,O,-representa the surface complex formed between a surface Mn(1II) center and adsorbed oxalate. It is assumed that the electron transfer in reaction 10is not reversible because the oxidized oxalate undergoes a fast decarboxylation reaction yielding one COz and one Con'- radical.44 The latter is a very reactive reductant that rapidly reduces a second surface Mn(II1). In this study, adsorption of oxalate onto the manganese(II1,IV) (hydrloxides could not be directly measured because of the rapid reduction of the (hydr)oxides. The plateau in the dissolution rate observed at high oxalate concentrations, however, supports the proposed mechanism in which the rate of reductive dissolution is directly proportional to the concentration of the surface complex. If the rate (in mol m-, min-l) could be normalized for the surface concentration, a rate constant (in min-l) could be obtained. It is, of course, this rate constant, rather than the rate, that should ideally be considered in evaluating the effect of parameters, such as pH, on the overall dissolution reaction. Dependence of the rate on the concentration of the surface complex demonstrates that a surface reaction, i.e. a reaction involving a surface species, must be the ratelimiting step in the reductive dissolution reaction. Oxalate adsorption may be excluded since a slow adsorption step should not produce the observed saturation effect. Effect of pH on the Dissolution Rates. The strong pH dependence of the rate of reductive dissolution of both y-MnOOH and P-Mn02 in the presence of oxalate does not correspond to any significant change in the thermo-
7
PH Figure 8. Summary plot showing the pH dependence of the rate of reductive dissolution of the solid manganese phases investigated in the presence of oxalate in the dark and in the light, with an initial oxalate concentration of 1 mmol L-l; A, y-MnOOH in the dark; A,y-MnOOH in the light; 0 , P-MnOz in the dark; 0,j3-MnOz in the light; m, N ~ M n l 4 0 ~ , . 9 Hin~ 0the dark.
Mn(I1). Second, when the morphology of the manganese(111)hydroxide particles (after exposure to oxalate at pH 4) was characterized by transmission electron microscopy,
(36) Morgan, J. J.;Sung, W.; Stone, A. T. In EnuironmentalInorganic Chemistry; Irgolic, K.; Martel, A. E., Eds.; VCH Publishers, Inc.: Deerfield Beach, FL, 1985. (37) Furrer, G.; Stumm, W. Ceochim. Cosmochim.Acta 1986,50,1847. (38) Zinder, B.; Furrer, G.; Stumm, W. Geochim. Cosmochim. Acta 1986,50, 1861. (39) Laha, S.; Luthy, R. Enuiron. Sci. Technol. 1990,24,363. (40) Stone, A. T.; Ulrich, H.-J. J.Colloid Interface Sci. 1989,132,509. Stumm, W. Langmuir 1991, 7, 809. (41) Suter, D.; Banwart, S.; (42) Sulzberger,B.;Suter, D.; Siffert, C.; Banwart, S.; Stumm, W. Mar. Chem. 1989, 2,23; 127. (43) Stumm, W.; Sulzberger, B.; Sinniger, J. Croat. Chem. Acta 1990, 63, 217. (44) Prasad, D. R.; Hoffman, M. Z. J . Chem. Soc., Faraday Trans. 2 1986,82, 2275.
Langmuir, Vol. 8, No. 1, 1992 101
Reductive Dissolution of MnOOH and MnOz by Oxalate
dynamic driving force for the reaction over the experimental pH range. For both (hydr)oxides, reaction with oxalate is very favorable even a t pH 6 (or above), particularly as the product C02 is continuously removed from the system. Several phenomena may plausibly contribute to the observed pH dependence. Protonation of the (hydr)oxide surface accelerates the nonreductive dissolution of (hydr)oxide minerals3'a8 and may, similarly, accelerate the detachment of the reduced surface manganese ions. Readsorption of Mn2+becomes more important with increasing pH40andmay, at higher pH, block the (hydr)oxidesurface, thus limiting adsorption of oxalate. And, the surface oxalate concentration may be expected to decrease with increasing pH, following the characteristic pattern for anion adsorption on (hydr)o~ides.4~ In order to examine this last pH effect, we calculated the pH dependence of oxalate adsorption using the intrinsic acidity constants of y-MnOOH and, because of lack of data, the intrinsic equilibrium constant of oxalate adsorption at the surface of goethite, a-FeOOH (Figure 9a). Note, that variations of 2-3 orders of magnitude of the latter constant do not change the general trend in adsorption as a function of pH (Figure 9b). Comparison of the experimentally determined rate of reductive dissolution of y-MnOOH as a function of pH with the speciation calculation for oxalate adsorption (Figures 8 and 9) indicates that the strong decrease in the rate of reductive dissolution with increasing pH parallels the pH dependence of the adsorption of oxalate at the y-MnOOH surface. The pH dependence of adsorption of weak acids and anions can, with the help of the ligand exchange model, be predicted from the acid-base equilibria of both the (hydrloxide and the anion.26 The pH,,, values of yMnOOH and of P-Mn02 are r6.2 and 2 < pH,,, < 4, re~pectively.~~ Thus, it is to be expected that the extent of oxalate adsorption is maximal at a higher pH for yMnOOH than for P-Mn02. This is consistent with the observation that thermal reductive dissolution of P-MnOz by oxalate does not occur to a measurable extent (within the time frame of our experiments) at pH > 5, while the rate of y-MnOOH dissolution remains measurable up to pH = 6 (Figure 8). This discussion does not exclude, however, the possible importance of Mn2+adsorption or proton catalysis of the detachment step. Further research is needed to address these issues. Facility of Thermal Reductive Dissolution and the Effect of Light. The oxidant efficiency and reactivity of manganese oxides is apparent from the standard potentials of the redox reactions Mn"'OOH(s)
+ 3 H + + e- s Mn2+(aq)+ 2H,O Eo = +1.50 V
Mn'"O,(s)
(12)
+ 4H+ + 2e- s Mn2+(aq)+ 2H20 Eo = +1.23 V
(13) which may be compared with the standard redox potential for the reaction 2CO,(g) + 2e- F-? C 2 0 t - Eo = -0.633 V (14) The facile thermal reduction observed on reaction of manganese(II1,IV) (hydrloxides with oxalate may be contrasted with the metastability of iron(II1) (hydr)oxides, such as goethite or hematite, in the presence of o ~ a l a t e . ~ In '!~~ part, this may reflect the larger thermodynamic driving (45) Schindler, P. W.; Stumm, W. In Aquatic Surface Chemistry; Stumm, W., Ed.; Wiley-Interscience: New York, 1987; Chapter 4. (46) McKenzie, R. M. Aust. J. Soil Res. 1981, 19, 41.
2
3
4
5
6
7
8
PH
-11
-
force for the reduction of Mn(1V) and Mn(II1) compared to Fe(II1). This can be seen, for example, by comparing the standard Gibbs free energy of the reductive dissolution reaction by oxalate of y-MnOOH (eq 4, AGO = -411.6 kJ mol-l) with the less favorable value of the corresponding reaction for goethite (a-FeOOH) (AGO = -276.4 kJ mol-'). The effect of light on the reductive dissolution by oxalate again shows a marked contrast between manganese(II1,IV) and iron(II1) (hydr)~xides.~l In the reduction of hematite by oxalate, formation of an electronically excited state enables the system to overcome the free energy of activation of electron transfer. In the case of manganese(II1,IV) (hydr)oxides, the energy barrier can be overcome by available thermal energy and light has only a minor effect on the reductive dissolution rate. A qualitative understanding of the thermal and photochemical reactivity of manganese(I11,IV) and iron(II1) (hydr)oxides may be obtained from frontier molecular
Xyla et al.
102 Langmuir, Vol. 8, No. 1, 1992 a R
4.H
ti
11
Q
R K R X
b PX
42
R
I
C
tetragonally distorted due to the Jahn-Teller effe~t.~E The difference in the reactivity of manganese(II1,IV) and iron(111) (hydr)oxides toward thermal reductive dissolution by oxalate may be rationalized as follows (compare Figure 10): In the thermal reductive dissolution of y-MnOOH and @-MnOzby oxalate, one and two electrons, respectively, from the u ligand orbital are transferred to the empty eg(u*)orbitals of the metals. Thus, electron transfer occurs along the bonding axes of the manganese(II1,IV) oxalato surface complex. Thermal reductive dissolution of iron(111)(hydrloxides, on the other hand, involves the transfer of an electron from a ?r ligand orbital to the tzg(7r) metal orbital, that is, an electron transfer that does not occur along the bonding axes of the surface complex, and thus involves a higher free energy of activation. This may explain why reductive dissolution of iron(II1) (hydr)oxides does not occur thermally in the presence of 0xalate.3~ In the photochemicalreductive dissolution of y-MnOOH and P-MnOz, various electronic transitions, due to absorption of light, may be involved. A d d transition (from the tzgto egorbital) would leave one A metal orbital vacant and available for electron transfer from a A ligand orbital (analogous to the electron transfer indicated in Figure 10for the thermal reaction of iron(II1) (hydr)oxide). (Note that, in contrast, a d d transition is not involved in the photochemical reductive dissolution of iron(II1) ( h y d r ) o ~ i d e s ~since ~ ~ ~such 9 ~ 5a~transition would not lower the free energy of activation of the electron transfer; compare Figure 10.) Alternatively, a ligand-to-metal charge-transfer transition of either or both the bulk manganese(II1,IV) (hydr)oxide and the surface complex may occur, that is, a transition from a ligand orbital, of the lattice oxygen or the surface oxalate, to a metal ligandfield orbital. The energetics of these electronic transitions vary significantly. The calculated energy difference between the 2 tzg and 3 eg orbitals is 3.33 eV (372 nm) for the (MnO@ cluster and 2.42 eV (512 nm) for the (Mn06)!+ cluster, whereas the MnlIIJV O-"charge-transfer transitions of these clusters occur in the UV region (below 300 nm).51 A ligand-to-metal charge-transfer transition within the surface oxalate complex can be expected to occur in the near-UV region, in analogy to the corresponding transition of the manganese(II1) oxalato solution comp l e ~ .In~ ~ our experiments, light below 350 nm was excluded and only a slight photochemical enhancement of the reductive dissolution was observed. If a greater enhancement of the reductive dissolution were to be observed with higher-energy light, this would demonstrate that a MnrrIJV O-" charge-transfer transition was involved in the photochemical reductive dissolution of yMnOOH and @-Mn02. Study of the wavelength dependence of the photochemical dissolution rate would also provide information on the role of a d d transition in the reductive dissolution of manganese(II1,IV) (hydr)oxides. The involvement of both the ligand-to-metal chargetransfer transition of the surface complex and the ligandfield excitation (i.e., a d d transition) in the photochemical reductive dissolution of y-MnOOH can be supported by comparison with the corresponding homo-
-
-
-
bg(4
'1
I
1
YZ XY
Figure 10. Molecular-orbital diagrams demonstrating the electron transfers involved in the thermal reduction of different metals. Splitting of the metal d orbitals in an octahedral ligand field (as a result of interaction with an oxygen donor atom of oxalate) is shown for (a) Mn(III), (b) Mn(IV), v d (c) Fe(II1). As noted in the text, the Mn(II1) geometry is subject to tetragonal distortion (not shown). Note, that thermal reduction of Fe(II1) by oxalate does not occur because of the too high free energy of activation.
orbital theory.47 The electronic configurations of the reacting metals are d3 for Mn(IV), d4 for Mn(III), and d5 for Fe(II1). The electron configurations of these metal ions in an octahedral ligand field are shown schematically in Figure 10 for the d electrons. The geometries of Mn(IV), in B-MnOp, and of Fe(III), in a-FeOOH, are regularly octahedral, while the Mn(II1) geometry, in y-MnOOH, is (47) Luther, G. W. I11 In Aquatic ChemicalKinetics: Reaction Rates of Processes in Natural Waters; Stumm, W., Ed.; Wiley-Interscience: New York, 1990; Chapter 6.
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(48) Wells, A. F. Structural Inorganic Chemistry, 5th ed.; Clarendon Press: Oxford, 1984. (49) Faust, B. C.; Hoffmann, M. R. Enuiron. Sci. Technol. 1986, 20, 943. (50) Litter, M. I.; Blesa, M. A. J. Colloid Interface Sci. 1988,125,679. (51) Sherman, D. M. Am. Mineral. 1984,69, 788. (52) Balzani, V.; Carassiti, V. Photochemistry of Coordination Compounds; Academic Press: London, 1970.
Reductive Dissolution of MnOOH and MnOz by Oxalate
Langmuir, Vol. 8, No. 1, 1992 103
metal center into the solution. The rates of the reductive geneous photoredox reactions. The photoredox decomdissolution are strongly pH dependent and much higher position of the manganese(II1) oxalato solution complex, at low pH values. This pH dependence parallels the pH M ~ ( O X ) ~occurs ~ - , through a ligand-to-metal chargedependence of the adsorption of oxalate on the surface, transfer transition with near-UV light and through a although other effects of protons on the kinetics of the ligand-field transition with visible light.52 However, these reductive dissolution of these manganese(II1,IV) (hydr)mechanisms are only pertinent to y-MnOOH if the surface oxides cannot be excluded. Comparison between the two oxalate complex (and not the bulk oxide) is the chro(hydr)oxide surfaces studied showed that the reactivity of mophore in the reductive dissolution. y-MnOOH toward reductive dissolution by oxalate is Implications for the (Photo)reactivity of Mangahigher than that of P-Mn02. The effect of light on the nese Oxides in Natural Waters. It has been found that reductive dissolution in the presence of oxalate was not for the reductive dissolution of manganese(1V) oxides in pronounced for either manganese (hydr)oxide. This is in the presence of fulvic acid, the rate and extent of to the results of studies where natural organics dissolution are markedly enhanced on i l l ~ m i n a t i o n . ~ ~ * ~ ~ 1contrast 53 were used as reductants, which suggests that light may act The fact that in the present study the effect of light was indirectly, in these cases, by producing secondary electron rather weak suggests that light might have an indirect donors. effect on the photodissolution on manganese(II1,IV) (hydrloxides, that is, by the secondary production of electron donors (e.g. H202) which are very effective in the Acknowledgment. We thank Rudolf Giovanoli (Unireduction of manganese(II1,IV) (hydr)oxides. The latter versity of Bern) for the characterization of the solids and has been shown recently by S t ~ r z e n e g g e r These . ~ ~ studies Beat Miller and Bettina Bartschat (EAWAG) for assisinclude two steps: (i) photochemical formation of hydrogen tance in the modeling. For helpful comments we thank peroxide and (ii) reductive dissolution of manganese(II1,James J. Morgan (CaliforniaInstitute of Technology),Alan IV) (hydrloxides in the presence of H202 in the dark. Stone (The Johns Hopkins University), and Patrick V. Brady (Southern Methodist University). This project is Conclusions part of the EEC Program European River Ocean Systems (EROS 2000) and has been financed by the Swiss ConThe reductive dissolution of y-MnOOH and P-Mn02 in federation, represented by the Federal Office for Education the presence of oxalate occurs thermally, at low pH values and Science, Federal Department (Ministry) of Home (from 4 to 6). The proposed mechanism includes the Affairs, Bern. formation of a precursor surface complex with oxalate followed by electron transfer and release of the reduced Registry No. y-Mn02H, 12025-99-9; P-MnOz, 1313-13-9; (53)Sunda, W. G.; Huntsman, S. A. Eos 1987,68, 1761. (54)Sturzenegger, V. T.Ph.D. Thesis, ETH Ziirich, No. 9004,1989.
Na2C20d,62-76-0;Mn(II), 7439-96-5;pyrolusite, 14854-26-3;manganite, 1310-98-1.
