Environ. Sci. Technol. 1984, 18, 617-624
Reduction and Dissolution of Manganese( I I I ) and Manganese( I V ) Oxides by Organics: 2. Survey of the Reactivity of Organics Alan T. Stone* and James J. Morgan W. M. Keck Laboratories of Environmental Engineering Science, California Institute of Technology, Pasadena, California 9 1125
Reduction and dissolution of manganese(II1, IV) oxide suspensions by 27 aromatic and nonaromatic compounds resembling natural organics were examined in order to understand the solubilization reaction in nature. At pH 7.2, M formate, fumarate, glycerol, lactate, malonate, phthalate, propanol, propiona?dehyde, propionate, and sorbitol did not dissolve appreciable amounts of oxide after 3 h of reaction. The following organics did dissolve manganese oxides under these conditions and are listed in order of decreasing reactivity: 3-methoxycatechol catechol 3,4-dihydroxybenzoic acid ascorbate > 4-nitrocatechol > thiosalicylate > hydroquinone > 2,5-dihydroxybenzoic acid > syringic acid > o-methoxyphenol > vanillic acid > orcinol 3,5-dihydroxybenzoic acid > resorcinol > oxalate pyruvate salicylate. Relative reactivities of organic substrates are discussed in terms of surface complex formation prior to electron transfer. Dissolution of manganese oxides by marine fulvic acid was enhanced by illumination, verifying that the reaction is photocatalyzed.
-
-
-
-- -
Introduction
Manganese oxide particles and crusts in natural waters can be reduced and dissolved by organic compounds, increasing the mobility of manganese and its availability to organisms. Dissolution by humic compounds, as well as by low molecular weight organics of similar structure, has been reported (1-3). Natural organics in marine surface waters have recently been shown to dissolve manganese oxides via a photocatalyzed reaction (4). A study was undertaken to understand the reductive dissolution process more fully and determine factors that influence its rate. Biological processes are not within the scope of this study. The first paper in this series (5) developed a rate law for the dissolution of manganese oxide suspensions by hydroquinone and examined inhibition by specifically adsorbing calcium and phosphate ions. Organic compounds having widely varying structures are present in natural waters, and their ability to reduce and dissolve manganese oxides may differ considerably. In order to predict rates of solubilization of manganese oxides, the reactivities and relative concentrations of organic compounds in each natural situation must be known. In addition, oxidation by manganese oxides may be an important degradative pathway for organic compounds in some environments. Manganese oxides have been shown to initiate coupling and polymerization reactions of phenolic compounds and may catalyze reactions with oxygen (6). Reactions of this kind may be important in the formation of humic material. Manganese oxides may also oxidize synthetic organics that are resistant to other decomposition reactions. Experiments are described in this paper in which manganese oxides are dissolved by a variety of simple organic *To whom correspondence should be addressed at the Department of Geography and Environmental Engineering, The Johns Hopkins University, Baltimore, MD 21218. 0013-936X/84/0918-0617$01.50/0
compounds having structures similar to organic material in natural waters. A marine fulvic acid sample is included in this study to allow comparison with natural organic material. Measurements of dissolution rate provide information about the reaction mechanism and a useful reactivity scale for different organic structural groups. These results are considered in light of the surface site-binding model for the dissolution reaction presented earlier (5). A number of important classes of organic compounds have been identified in aquatic systems. Metabolites and simple biological molecules have been identified in waters and sediments, including amino acids, carbohydrates, lipids, heterocyclic compounds, vitamins, phenols, and quinones (7,B). Monosaccharides and sugar acids (uronic acids) are produced in large amounts by marine algae and are readily identified in marine sediments (9, 10). Dissolved organics in pore waters are formed from in situ microbial reaction with insoluble organic detritus (11). Low molecular weight metabolites such as formic, acetic, l-butyric, and isobutyric acids have been measured at M in reducing pore concentrations exceeding 1.0 X waters and may make up as much as half of the total dissolved organic carbon in pore waters (11). Glycolic, lactic, oxalic, and succinic acids have also been identified in reducing sediments (12). Humic substances are the predominant organic compounds in most natural waters. Chemical degradation studies of freshwater humics indicate that humics contain a core structure of phenols and phenolic acids such as hydroxybenzoic acids, vanillic acid, syringic acid, 3,4-dihydroxybenozoic acid, 3,5-dihydroxybenzoic acid, resorcinol, and catechol (13). These aromatic groups are linked together by short, saturated aliphatic chains, possibly at three or more positions on the aromatic ring (14). Similar studies with marine humics reveal a predominantly aliphatic structure resembling cross-linked, autoxidized, polyunsaturated fatty acids (15). Freshwater and marine humics have such different structures that considerable differences in reactivity are likely. Organic Substrates
Fifteen aromatic and 12 nonaromatic compounds were chosen to represent the variety of structures and oxygencontaining functional groups found in natural organic material. Compounds studied, their source, and their pK, values are listed in Table I. All experiments were performed at pH 7.2, unless otherwise noted. With the exception of 4-nitrocatechol, aromatic hydroxy groups of phenolic substrates are fully protonated at this pH. One hydroxy group of 4-nitrocatechol is protonated and the other deprotonated at pH 7.2. Standard potentials for organics that can be oxidized reversibly in acidic solution are listed in Table 11. The manganese(II1, IV) oxide phase used in the experiments contained feitknechtite, manganite, and possibly birnessite (21); standard potentials for these phases are also given in Table 11. AG, the Gibbs free energy of the dissolution reaction, can be calculated by using Eo values provided that organic substrates are oxidized to the same products
0 1984 American Chemical Society
Environ. Sci. Technol., Vol. 18, No. 8, 1984
617
-__
--_____.---
--
___- __
..-
__
Table T. Organic Compounds Included in the St,udy, Their pK, Values, and Their Source rompound
PK,,
PKa2
ascorhic acid* catecholh 2,s-dihydroxybenzoic" acid (2,5-DIOH) 3,4-dihydroxybenzoica acid (3.4-DTOH) 3,5-dihydroxybenzoicnacid (3,5.DTOH) EDTA' formic acide fumaric acidd
4.37 9.34 2.97 4.26 8.84 1.994 3.53 3.02 14.9 10.03. 3.57 2.62 9,20
11.34 12.6 10.21 8.64 9.00 2.674
glyceroln
hydroquinoneb lactic acid" malonic acidh 3-meth~xycatechol~ 2-methoxyphenol" 4-nitrocatechol" orcinol" oxalic acid' pht,halic acidf propanoic acid" pyruvic acid" l-propanola propionaldehyde" resorcinolh salicyclic acidb sorbitolb syringic acid" thiosalicyclic acid" vanillic acid"
ionic strength, M, and temp, "C
PK83
3.0, 25 0.1, 25 0.1, 31 0.25, 25 0.2, 25 0.1, 20 1.0, 20 0.04, 18
13.90 13.33 10.54
6.17
4.45 11.66 5.30 11.99
9.90
6.70 9.36 1.26 2.76 4.52 2.35
10.85 1.1.66 3.82 4.92
9.30 2.80 13.00 4.34 3.6 4.51
11.06 13.6 9.49 8.2 9.39
ref
16 16 16 16 16 18
16 19
0.1, 13 0.2, 25 0.1, 25 0.1, 25 0.1, 25 0.1, 20 0.1, 20 0.1, 25 0.1, 22 0.3, 25 0.5, 25
17 17 16 16 20 17 16 17 16 16 16 16
0.1, 20-25 0.1, 25 ?, 60 ?, 25 0.1, 25 ?, 25
17 16 17 I6 16 17
"Aldrich. b8igma. J. T. Baker. 'Mallinckrodt. /Matheson. - - .. .. . __"MCR. -. . . .- . . . - .. __ --..~.. ..- ___ . . _.._I . .. ..~ -. .. . __ substrates listed in Table I1 under the experimental conTable 11. St.andard Potentials of Selected Half-Reactionsn - ..
~
~
reaction 2CO,(g) + 2H+ + 2e- = HzC20, (oxalic acid) COz(g) 2H+ + 2e- = HCOOH (formic acid) CH,COCOOH (pyruvic acid) + 2H+ + 2e- = CH,CH(OH)COOH (lactic acid) dehydroascorbate + 2H+ + 2e- = ascorbate Q (quinone) + 2H+ + 2e- = QH, (hydroquinone) Q (quinone) + 2H+ + 2e- = QH, (2,5-DIOH) Q (quinone) + 2H+ + 2e- = QH2 (catechol) Q (quinone) + 2H+ + 2e- = QH2 (3,4-DIOH) Q (quinone) + 2H+ + 2e- = QH, (4-nitrocatechol) P-MnOOH(s) + 3H+ e- = Mn2++ 2Hz0 r-MnOOH(s) + 3H+ + e- = Mn2++ 2HzO 1/26-Mn0,(s) + 2H+ + e- = la2Mn2++ HzO Fe(OH),(s) + 3H+ + e- =I Fe + 3H20
V ref 4,4922 -0.196 22 +0.20 22 EO,
+
+0.40 23 +0'699 24
+
+o,77 +0,792 +0.883 +0.95
25 24 23 25
+1.65 +1.50
26 27 27 28
+0.95
+
P-MnOOH(s) = feitknechtite; r-MnOOH(s) = manganite; 6MnO,(s) = birnessite; 2,5-DIOH = 2,5-dihydroxybenzoic acid; 3,4DIOH = 3,4-dihydroxybenzoic _- acid.._ - -
given in the half-reactions. Oxidation of hydroquinone to p-benzoquinone by MnOOH(s) is described by reaction 1. MnOOH(s) + 1/2QH2+ 2H+ = Mn2+ + l/*Q + 2H20 (1) A G is given by the following relation: =
-F@MnOOH
(2)
- Ehydroquinone)
Equation 2 can be expanded by using the Nernst equation: AG =
-F(EoMnOOH
- E'hydroquinone)
RT In
f.
(
)
[Mn2f][Q]1/2 - -
[H']2[QH2]1/2
Environ. Sci. Technol., Vol. 18, No. 8 , 1984
(-72):
-
R--SS--R + 2Fe(CN)64-+ 2H+ (4) Phenols such as salicyclic acid are considerably more resistant to oxidation than dihydroxybenzenes and methoxyphenols, although phenoxy radicals can also be formed. Phthalic acid, containing no hydroxy or methoxy substituents, i s not expected to react at all. Oxidations by Mn3+(aq) and Mn(1II) complexes in strongly acidic solutions have been extensively studied. Oxidation of organics by solid manganese oxides may involve binding of organics directly to manganese atoms on the oxide surface and may therefore resemble oxidations by dissolved Mn(II1) species (21). 2R-SH
+ 2Fe(CN),"-
Experimental Methods
(3)
AG is negative and the reaction favorable for dissolution of feitknechtite, manganite, or birnessite by the organic 618
ditions employed. Oxidation of organics by manganese oxide surfaces may form the same oxidized products formed by similar inorganic oxidants. Catechols and hydroquinones are oxidized reversibly to o- and p-benzoquinones by some oxidants and irreversibly to coupled and ring-cleavage products by others. Resorcinols cannot be oxidized to benzoquinones but can be oxidized to coupled and ring-cleavage products via radical reactions (29). Suitable oxidants such as Tl('II1) can oxidize o- and p-methoxyphenols to benzoquinones through loss of the methoxy group as methanol (30). Vanillic and syringic acids can both be oxidized to obenzoquinones in this way. One-equivalent oxidation of thiosalicylate (using ferricyanide) forms a dimer linked through a disulfide bond
All stock solutions prepared from solid reagents were filtered with 0.2-pm pore diameter membrane filters (Nucleopore Corp.) prior to use. Stock solutions of organic substrates were used within 24 h to minimize oxidation by oxygen. In most cases, solutions were made slightly acidic
0 Time (Minutes)
Figure 1. Dissolution of manganese oxide suspensions under set I (low reductant concentration) conditions: (0)methoxycatechol; (A) $4DIOH; (0)catechol; ( 0 )nitrocatechol; (0)hydroquinone; (U)2,5-DIOH.
to retard autooxidation. Marine fulvic acid was provided by G. R. Harvey of the Atlantic Oceanographic and Meteorological Laboratories (NOM). The sample had been adsorbed on activated carbon from seawater passed through an XAD-2 resin. Characterization of similar material has been described by Harvey et al. (15). Dissolution experiments followed procedures outlined previously ( 5 ) . Manganese oxide suspension N(9), described in ref 5 and 21, was used. This is a mixed phase containing feitknechtite and lesser amounts of manganite and an Mn(1V) phase (21). I t has a composition characterized by MnO, fi4 and a pH,, (zero proton condition) of 6.9. The pH of reaction solutions was maintained at 7.2 with a bicarbonate buffer (lo-, M NaHCO,, 1% COJ, and ionic strength maintained at 5.0 X lom2 M with NaNO, (or by NaClO,, as in the spectral analysis experiments). Reductants that react quickly with manganese oxides were studied under set I conditions: 1.77 X M reductant mol/L manganese oxide suspension and 3.54 X [(MnO,Io). Set I conditions are approximately stoichiometric for two-equivalent reductants. Set I1 conditions, having elevated reductant concentrations, were used with less reactive substrates: 1.00 x lo-, M reductant and 2.83 X mol/L [MnOJo. Dissolved manganese was measured in aliquots of reaction solution filtered with 0.2 wm pore diameter by 25mm membrane filters (Nucleopore Corp.) by using the techniques described ref in 5. Measurements of UV absorbance were made immediately after filtration. Results Relative Rates of Dissolution. Experiments performed with different substrates are listed in Table 111. [Mn2+Idi,,is the amount of dissolved manganese determined by the filtration technique outlined previously (5). In blank experiments containing no organic reductant, [MnZ+ldiRs was less than 3% of the total manganese added (at pH 7.2). In experiments marked NR (no reaction), [Mn2+Idi,measured 3 h after the addition of substrate was no larger than in the blank experiments. No reaction was observed when formate, fumarate, glycerol, lactate, malonate, phthalate, propanol, propionaldehyde, propionate, or sorbitol was used as substrate under set I1 (excess substrate) conditions. The amount of oxide dissolved by pyruvate, oxalate, or salicylate was small under set I1 conditions, barely above levels in the blank solution. An experiment in which the pyruvate concentration was increased to M confirmed that pyruvate was indeed dissolving the oxide. Less than 10% of the oxide was dissolved. Ascorbate was the only nonaromatic substrate studied that dissolved the oxide suspension under set I (low sub-
50
100
150
200
250
Time (Minutes)
Flgure 2. Dissolution of manganese oxide suspensions under set I conditions: ( 0 )ascorbate; (0)thiosalicylate; (0)syringic acid.
/
I
Time (Minutes)
Figure 3. Dissolution of manganese oxide suspensions under set I1 (excess substrate) conditions: (e)o-methoxyphenol; (B)vanillate; (0) 3,5-DIOH; (0) orcinol; (A) resorcinol.
strate concentration) conditions. In fact, ascorbate dissolved the suspension as quickly as any of the substrates tested. Results from dissolution experiments under set I conditions are shown in Figures 1 and 2. Ascorbate, 3methoxycatechol, 3,4-DIOH, and catechol dissolved the suspension at approximately the same rate. 4-Nitrocatechol, with a strong electron-withdrawing substituent group, dissolved the oxide at a rate one-fourth that of other catechols. Hydroquinone and 2,5-DIOH, both p-dihydroxybenzenes, dissolved the oxide more slowly than any of the catechols (o-dihydroxybenzenes). With sufficient time, 0-and p-dihydroxybenzenes dissolved all the oxide under set I conditions. These compounds are therefore two or more equivalent reductants under these conditions. Thiosalicylate initially reacted with the oxide at an appreciable rate but dissolved only 1.2 X M of the oxide. This point was reached within 2 h of reaction. Since there are 1.3 equiv of oxidant/mol of manganese oxide and the thiosalicylate concentration was 1.77 X mol/L, then 0.87 equiv of manganese oxide was reduced per mol of thiosalicylate. Syringic acid, a dimethoxyphenol, reacted very slowly under set I conditions (Figure 2). The reactivities of other monophenols and of resorcinols were still lower, and no appreciable dissolution was measured with set I conditions. Increasing the substrate concentration 50-fold to set I1 conditions made it possible to measure rates of dissolution. Vanillate and o-methoxyphenol reacted more quickly than the resorcinols, as shown in Figure 3. In order to compare the reactivities of the substrates, rate cnnstants independent of substrate concentration and suspension loading need to be calculated. For our purposes Environ. Sci. Technol., Vol. 18,
No. 8, 1984 619
k, (liters/mole.sec., a t paH 7 . 2 )
@
Catechols &OH OH
3x10' 2x10' 2x10' 2X101
3-Methoxycatechol Catechol 3,4-Dihydmxybenzoic Acid Ascorbate
5.33~10'
4-Ni trocatechol
COOH catechol 3,4-dihydroxybenzoic acid @ Hydroquinones
4-nitrocotechol
OH
bH 2,s-dihydroxybenzoic acid
6H hydroquinone
Syringic Acid
k0,
3-methoxycatechol
Q
1 . 7 2 ~ 1 0 ~ 2,5-Oi hydroxybenzoic Acid
&OH
OCH. a
OH
2.68~10~ Thiosalicylate 2.33~10' Hydroquinone
5.75~10-I
&OH
$OH
@ Resorcinols
3.25x10-'
V a n i l l l c Acid
1.32x10-~ 1.24~10"
c I
Orcinol 3 . 5 4 1 hydroxybenzoic Acid
4.90~10-~ Resorcinol
&lae3:
6 HobcooH HoA
o-Methoxyphenol
8.63x10-'
Pyruvate, Oxalate, Salicylate
Figure 4. Apparent second-order rate constants for dissolution of manganese oxide suspensions. These are rough values presented for the sake of comparison only.
HOresorcinol
3,5-di hydroxy CH3 toluene
3,5-di hydroxy benzoic ocid
@ Methoxy
(orcinol)
Aromatics
F
H
0-methoxyphenol (guaiacol)
S
H
HOOC Vanillic Acid
@ Mono-Substituted
Z
t
H
HOOC OCHJ OH Syringic Acid 3-methoxycatechol
Benzoic Acids
COOH
it is assumed that the orders with respect to reductant concentration and manganese oxide loading are the same as observed in the reaction with hydroquinone (5). The following rate law is used for all substrates at constant p H
E
H
COOH
Salicylic Acid
@
Thiosalicylic Acid
Non Aromatics
L-ascorbic acid O=C
I 1 _I"
HO-C
Under set I conditions, the reductant is depleted along with the oxide suspension. In order to calculate k, for reactions under set I conditions, the stoichiometry of the reaction must be known. It is assumed that all organics are twoequivalent reductants. This assumption is clearly incorrect for thiosalicylate and possibly for other substrates. Rate constants are therefore rough values calculated for the sake of comparison only. Rate constants were calculated from the experimental data by using the method outlined in a preceding paper (5). Figure 4 gives apparent second-order rate constants (k,) for substrates that reduced the oxide, and Figure 5 shows their structures. Rate constants calculated for this group of substrates differ by more than 3 orders of magnitude. On the basis of these results, the following generalities can be made concerning structure/reactivity relationships: (i) for dihydroxybenzenes, the order of reactivity is ortho (catechols) > para (hydroquinones) >> meta(resorcinols), (ii) for methoxy-substituted phenols, the order of reactivity is dimethoxyphenols > monomethoxyphenols >> phenols, (iii) strong electron-withdrawing substituents (such as -NO2) lower substrate reactivity, and (iv) substitution by a carboxy group may either promote or inhibit the reaction. Spectral Analysis. UV and visible spectra were recorded during the oxidation of the nine most reactive organics, and the results are summarized here. The reader interested in more detail should refer to ref 21. Set I condition (low substrate concentration) were used, and aliquots were filtered prior to placing in the spectrophotometer to remove oxide particles. 620
Environ. Sci. Technoi., Vol. 18, No. 8, 1984
HO-!
HO-CH i CH,OH pyruvic ocid CH,COCOOH Oxalic acid HOOC-COOH
Figure 5. Structures of organic compounds included in the study that reduce and dissolve manganese oxides (under conditions described in the text).
1
2.75F
2.50
I
i
I
i
I
1-1
2 0 MINUTES
W a v e i e n g t h (nm )
Figure 6. Spectra of filtered reaction solution during oxidation of hydroquinone by manganese oxide suspension. Conditions: 3.5 X M [MnO,],, 1.77 X M hydroquinone, paH 7.25, 5.00 X lo-' M NaCIO,, lo-, M NaHCO,, 1 % CO,, 25 OC, 10-cm cell.
Oxidation of hydroquinone by manganese oxide suspension forms p-benzoquinone. As Figure 6 illustrates,
Table 111. Dissolution Experiments with Selected Organic Substrates' trial
PBH
GA GB GC GD GE GF GG GH GI GJ GK GL
7.19 7.19 7.19 7.20 7.4 7.19 7.19 7.68 7.73 7.70 7.16 7.17 7.17 7.24 7.25 7.26 7.20 7.19 7.22 7.13 7.11 6.5-7.6 7.15 7.14 7.15 7.29 7.30 7.28 7.17 7.17 7.26 7.20 7.21 6.9 7.65-7.74 6.90-7.11 6.35 6.50
GM GN GO GP GQ GR GS GT GU GV GW GX GY GZ XA XB
xc
XD XE XF XG XH XI AB AI AJ
[MnO,l, M 2.82 X 2.88 x 2.88 X 2.88 X 2.88 x 2.88 x 2.88 X 3.43 x 3.43 x 3.43 x 3.54 x 3.54 x 3.54 x 3.54 x 3.54 x 3.54 x 3.54 x 3.54 x 3.54 x 2.83 x 2.83 X 2.83 x 2.83 x 2.83 X 2.83 x 3.54 x 3.54 x 3.54 x 2.63 x 2.63 X 3.29 X 2.64 x 2.64 X 2.64 x 2.85 X 2.81 x 2.89 x 2.89 X
10-5 10-5 10-5 10-5 10-6 10-6 10-5
10-6 10-6 10" 10-5 10-6 10-6 10-6 10" 10-5 10" 10-5 10-5 10" 10-5 10" 10-5 10-5
10-5 lo6 10-5
10-5 10-5 10-5
reductant
results
M propanol M propionaldehyde M propionic acid 10" M glycerol M lactate 10" M pyruvate M malonate M hydroquinone M catechol lo6 M resorcinol M 3,5-DIOH M 2,5-DIOH M 3,4-DIOH M orcinol M hydroquinone M catechol M syringic acid M o-methoxyphenol loF5M 3-methoxycatechol 10" M salicylate M vanillate M fumarate M orcinol 1.00 X M 3,5-DIOH 1.00 X M resorcinol 1.77 X M 4-nitrocatechol 1.77 X M ascorbate M thiosalicylate 1.77 X 1.00 X M o-methoxyphenol M oxalate 1.00 X 1.69 X 10" M formate 9.90 X lo4 M formate M pyruvate 1.00 X M 2,5-DIOH 2.63 X 199 mg/L fulvic acid 2.00 x 10-3M EDTA 201 mg/L fulvic acid blank
NR NR NR NR NR 5.6% dissolved NR Rxn Rxn NR NR Rxn Rxn NR Rxn Rxn Rxn 2,5-DIOH > catechol > 3,4-DIOH > 4-nitrocatechol. Relative rates of reaction with manganese oxides, however, follow a different order: catechol 3,4-DIOH > 4nitrocatechol > hydroquinone > 2,5-DIOH. Catechols react more quickly than hydroquinone despite less favorable free energies of reaction. Oxidation of catechols by manganese oxides did not form o-benzoquinones as assumed, however. E" values of catechols may not be representative of actual differences in free energies of reaction. o-Hydroxy substituents of catechols can form bidentate, inner-sphere complexes with dissolved metal ions, whereas p-hydroxy substituents of hydroquinones cannot. Chelation enhances the stability of catechol complexes relative to those of hydroquinones. Catechols have been shown to adsorb strongly to iron oxide (38)and aluminum oxide (39) surfaces. Evidence indicates inner-sphere coordination and bidentate surface complex formation with catechol (39). If a bonded mechanism is correct, increased surface coverage by catechols allows them to dissolve the oxide more quickly than hydroquinones, provided that rates of electron transfer within surface complexes are approximatelyequal. Rates of surface complex formation and electron transfer must be distinguished before either a bonded or nonbonded mechanism can be assigned to this reaction. It should also be pointed out that some organic compounds may react via a bonded mechanism, while others react via a nonbonded mechanism.
-
Summary Two of the 15 aromatic substrates and 8 of the 12 aliphatic substrates did not dissolve appreciable amounts of manganese oxide under the conditions studied. Saturated alcohols, aldehydes, ketones, and carboxylic acids showed no reactivity, except for pyruvic and oxalic acids. CateEnviron. Sci. Technol., Vol. 18, No. 8, 1984 623
chols, hydroquinones, methoxyphenols, and resorcinols, as well as ascorbate, reduced and dissolved manganese oxide suspension at appreciable rates. With the exception of ascorbate, molecules such as these have been shown to form the core structure of humic substances. Humic substances should therefore dissolve manganese oxides under natural conditions. Differences in the overall free energy of reaction do not correlate well with differences in relative reactivity. Factors influencing the formation of surface complexes and electron transfer at the surface appear most important. Electron-withdrawingsubstituents on aromatic substrates lower the reaction rate with manganese oxides, while electron-donating groups increase the reaction rate. Photoreduction of manganese oxides by marine humic substances, first observed by Sunda et al. (4), has been confirmed. Dissolution by marine fulvic acid was faster at lower pH values, most likely caused by greater adsorption of fulvic acid on the oxide surface. Literature Cited Hem, J. D. Geol. Suru. Water-Supply Pap. (U.S.) 1965, NO.1667-0. Baker, W. E. Geochim. Cosmochim. Acta 1973, 37, 269. Guy, R. E.; Chakrabarti, C. L., Can. J. Chem. 1976,54,2600. Sunda, W. G.: Huntsman, S. A.; Harvey, G. R. Nature (London)1983, 301, 234. Stone, A. T.; Morgan, J. J. Enuiron. Sci. Technol. 1984,18, 450. Larson, R. A.; Hufnal, J. M. Limnol. Oceanogr. 1980,25, 505. Degens, E. T. “Geochemistry of Sediments”; Prentice-Halk Englewood Cliffs, NJ, 1965. Cranwell, P. A. In “Environmental Chemistry”; Eglinton, G., Ed.; The Chemical Society: London, 1975; Vol. 1. Mopper, K.; Larson, K. Geochim. Cosmochim.Acta 1978, 42, 153. Mopper, K.; Dawson, R.; Liebezeit, G.; Ittekkot, V. Mar. Chem. 1980,10, 55. Barcelona, M. J. Geochim. Cosmochim.Acta 1980,44,1977. Peltzer, E. T.; Bada, J. C. Geochim. Cosmochim.Acta 1981, 45, 1847. Norwood, D. L.; Johnson, J. D.; Christman, R. F.; Hass, J. R.; Bobenrieth, M. J. Enuiron. Sci. Technol. 1980,14,87. Liao, W.; Christman, R.; Johnson, J. D.; Millington, D. S.; Hass, J. R. Enuiron. Sci. Technol. 1982, 16, 403. Harvey, G. R.; Boran, D. A.; Chesal, L. A.; Tokar, J. M. Mar. Chem. 1983,12, 119.
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Received for reuiew September 2,1983. Accepted March 14,1984.
