Environ. Sci. Technol. 1900, 22,62-70
Hausmannite (Mn,O,) Solution
Conversion to Manganite (y-MnOOH) in Dilute Oxalate
Carol J. Llnd Water Resources Division, U.S. Geological Survey, Menlo Park, California 94025
rn Oxalic acid retards the alteration of Mn304 to
y-
MnOOH during aging a t pH 7.4 f 0.2 in well-aerated, abiotic suspensions that contain 4.4 X M total Mn. M oxalate and greater, about In solutions of 1.25 X 15% of the initial Mn304altered to y-MnOOH by day 10, and in solutions of 6.7 X lo4 M oxalate, about 45% altered to y-MnOOH by day 67. Although precipitation continued through day 365, the degree of conversion remained the same as at day 10 and day 67, respectively. In oxalate-free suspensions, the conversion was about 80% complete by day 67 and 100% by day 109. Oxalate complexed most of the dissolved divalent Mn, lowered the free Mn(I1) and MnS02 concentrations, but increased the total dissolved Mn. Steric hindrance of surface reactions by a suggested manganese oxalate layer on the Mn304surface may explain the blockage of the oxidation cycle. Introduction The complexing capacity of inorganic anions has been shown to influence the Mn oxidation pathway and, consequently, the mineral form of particulate Mn (1). Because natural aqueous organic matter is intimately associated with trace elements (such as Mn) and their reactions, possible changes in the Mn oxidation pathway caused by the presence of a complexing organic anion also are of interest (2-4). This paper examines the alteration of the Mn oxidation pathway by an organic anion, oxalate; suggests a reaction mechanism for this alteration; and discusses some influences these alterations can have in natural aqueous settings. In many natural water settings, Mn exists primarily in particulate form as shown by the following: river particulates, 1050 gg/g; soils, 1000 gg/g; deep sea clays, 6000 gg/g; oceanic suspended matter, euphotic zone, 529 pg/g, as opposed to dissolved Mn in river water, 8.2 pg/L, and in ocean water, 0.2 kg/L (4-6). The ratio of particulate Mn to dissolved Mn is often in a state of flux ( 4 , 7). Mn in some estuaries may be a conservative element and in others precipitated as oxides and hydroxides, as particles or coatings, or adsorbed as Mn(I1) on organic materials (4). Sometimes particulates may dominate a harbor during flood events while dissolved Mn may dominate during normal flow conditions (8). In surface waters and estuaries, Mn(I1) may accumulate in the hypolimnion because of seasonally dependent Mn(I1) dispersion from reduced sediments into the overlying water column (7, 9, 10). Freshly mobilized Mn(I1) diffuses away from its source and when favorable conditions are encountered reprecipitates (11). The dissolved Mn concentrations in the epilimnion are low caused by rapid oxidative removal during turnover (9). Descending organic-rich detritus and PO, changes cause movements in the oxic/anoxic boundary (11). The oxidation state of Mn in these particulates usually is between +2 and +4. The identification of the actual Mn mineral form is obscured by the minute size of Mn precipitates and the multitude of other mineral and biological forms and organic materials with which they coexist. Thermodynamically, the physical and chemical solution parameters of oxic natural waters generally predict a form of MnO,, or an oxide with a slightly lower oxidation state. 62
Environ. Sci. Technol., Vol. 22, No. 1, 1988
In acidic to slightly alkaline anoxic waters, these parameters predict increased Mn(I1) concentrations and in more alkaline anoxic waters predict less oxidized Mn solids (12). Some stability relations of the Mn-0,-H,O systems were studied by Bricker (13), reaction products of Mn-bearing waters were described by Lind et al. (14), and rates of reductive dissolution of Mn are discussed by Stone (11). Besides conditions that favor direct oxidation to MnOz, there is a range of solution parameters that favor the formation of metastable manganite (y-MnOOH) as an intermediate along the oxidation pathway from Mn(I1) to Mn02. Under these conditions the mechanism governing transformation to MnO, may become more difficult to maintain and because y-MnOOH is oxidized slowly, yMnOOH may persist metastably. (The persistence of y-MnOOH will be illustrated by some of the data presented in this paper.) Under near natural conditions and in the absence of other solute or suspended matter, it has been suggested that the rate-determining step in yMnOOH formation is the transformation of the rapidly formed initial products hausmannite (Mn304)and feitknechtite (P-MnOOH) to manganite (y-MnOOH) (15). Therefore, conditions regulating y-MnOOH formation would influence the fate of aqueous Mn. Meaningful application of reaction kinetics to manganese chemistry in natural aqueous settings is a very complex chore because of the multitude of interactions occurring simultaneously, each at their own rate. Pankow and Morgan (16) discuss types, rates, and reversibilities of aqueous Mn reactions and how these factors control the existing state of Mn. These factors are influenced by pH, Eh, ionic strength, chemical composition of the water and of the solids the water contacts, biological activity, and the amount and the character of particulate surfaces available (6). Removal of trace metals from surface water appears to involve both physical and biological mechanisms. The primary trace metal concentration control mechanism seems to be physical and chemical adsorption onto biologically produced particulate matter (4). Dormant spores of marine Bacillus have been shown to significantly enhance the Mn oxidation rates at pH values below which Mn(I1) autoxidizes (In,and the action of bacterial poisons have been shown to inhibit the binding of Mn by bacterial catalysts (18). In the latter case, Mn was being oxidized and not simply bound. In deep waters, bacterial decomposition of particles releases trace metals back into the water column. In sediments, microorganisms may contribute to the reductive dissolution of manganese oxides by using manganese oxides as a terminal electron acceptor or through the production and excretion of reductants and other molecules that perturb the chemical conditions in sediments (11). In the oceanic water column, the form of aqueous Mn (dissolved and particulate) appears to primarily depend on kind, production rate, health, and age of the biota. These in turn depend on change in water chemistry caused by advection and mixing because of storms and turbulent events. During the late stages of an algal bloom, there may be high production of organic exudates with metal complexing ability. Control of most
Not subject to U S . Copyright. Published 1987 by the American Chemical Society
trace metals in surface water is related to local primary productivity (4). This paper illustrates how a biologically produced chemical can abiotically contribute to aqueous Mn chemistry. Several variables determine the Mn oxidation pathway. One variable is the nature of the major anions present (I, 19). If these anions are weakly complexing or noncomplexing (that is, chloride, nitrate, or perchlorate), the initial product at 0 "C is primarily p-MnOOH. If the major anion is sulfate, a Mn complexer, the initial products primarily are a-and y-MnOOH. If any of the above are the major anions present at 25 OC, the primary initial product is Mn304(the Mna04later alters to y-MnOOH). In the 0 "C sulfate medium, the MnS02 ion pair may inhibit development of the (Mn(OH)2), hexagonal platelet (the precursor of p-MnOOH) and thus cause the initial precipitates to be a-and y-MnOOH. Other ion pairs, such as Mn with organic ligands, will likely also alter the reaction pathway. Natural aqueous organic matter is intrinsically involved in the reactions that determine the fate of dissolved metals in natural water systems. These organic substances complex trace metals and mediate their adsorption and chemical reactions (20-29). Adsorbed organic material can mask properties of the underlying solid and present a surface with different physical and chemical properties (30-32). Natural organic substances are capable of reducing and dissolving manganese oxides, precipitating Mn, stabilizing the solubility of Mn as a complex, and inhibiting oxidation of Mn(I1) to Mn02 (2,3,20,33-36). Light has been reported to enhance manganese oxide reduction and dissolution by organic matter (33,36). Conversely, manganese oxides oxidize and polymerize organic matter (33, 37). Oxalic acid, the most simple form of dicarboxylic acid, was chosen to model organic matter-Mn interactions. Compared with many other naturally occurring organic compounds tested, oxalic acid ranked intermediate with regard to manganese(II1,IV)oxide dissolution rate (2,33). Oxalic acid has been measured in many natural water settings such as seawater (38),reducing sea sediments (39), and air (40) and at concentrations as high as 0.3 ppm in precipitation (40). Oxalate has been identified as a manganese oxalate precipitate in and beneath lichens growing on Mn ores (34). Oxalic acid is produced among the marlic acid oxidation products and by enzymatic oxidation of glycolic acid. [Malic acid is an important root exudate constituent of several plant species, and glycolic acid comprises a large fraction of algal exudate (35, 41-44).] Oxalate comprises from less than 0.1 to 5% of the dry weight of several plant species that are potential detrital imputs to freshwater and marine sediments. When wet, these sediments have been shown to contain 0.1-0.7 mmol of oxalate (45). Thus, besides using this acid as a model organic compound, oxalate concentrations are sufficiently high in many natural settings for the manganese-oxalate reactions themselves to exert an influence on the fate of aqueous Mn. Another reason for choosing oxalic acid is the Mnbinding capacity of this acid. Electron spin resonauce studies (ESR)have shown that organic matter must have more than one carboxyl site available to form tightly held, inner-sphere complexes with Mn, and when soil pH is high enough to make sufficient carboxylate sites available, inner-sphere Mn complexes can form with solid soil organics (20). Comparison of the ESR results and the properties of the simple acids considered in that study with properties of oxalic acid suggests that at pH values above neutrality the two carboxyl groups of oxalic acid would likely form
an inner-sphere chelate, MnC202, with Mn(I1). Oxalate introduction into a suspension of Mn,04 illustrates how this acid can influence the dissolved and solid form of Mn and should serve as a model for similar, and in some cases even more intense, influences of natural organic matter. Ultimately, these influences extend to factors that relate to Mn availability, distribution, and transport. Soluble, organically complexed Mn can aid or prevent biological absorption of Mn depending on the complex's ability to pass through the cell wall. Mn complexes, whether soluble, precipitate, or particulate coating, can be a reservoir that regulates Mn availability with regard to biological nutrition and toxicity (28, 34, 46, 47).
Experimental Methods Mn304Preparation. A suspension of Mn304 was prepared by adding 0.118 N NaOH to a stirred 0.01 M MnS04solution that was maintained at 25.0 f 0.1 OC while being flushed with C02-free air. A Radiometer pH-Stat controlled the microdrop addition of NaOH at a rate that kept the pH at 8.5 f 0.2. The base addition was continued until approximately 2 mol of base had been added for every 1mol of Mn. All chemicals used were of reagent grade or better. When the reaction rate became so slow that the base addition had essentially stopped, the reaction was considered complete, and no more base was added. The resulting suspended solid was in a 0.01 M Na2S04solution. The flushing of the space above the suspension with C02-free air and the stirring were continued. About 24 h after the initial NaOH addition, the pH had dropped to 7.82. Mn304Suspensions Containing Differing Oxalate Concentrations. A series of diluted suspensions was prepared by mixing 140-mL portions of the 24-hour-old suspension with 140-mL volumes of 0.01 M Na2S04containing differing sodium oxalate (Na2C204)concentrations. As the pH value of the oxalate-free, diluted suspension was 7.59 f 0.01, the pH values of the oxalate containing 0.01 M Na2S04solutions were each adjusted to 7.60 f 0.01 prior to mixing with the 140-mL aliquots of undiluted Mn304 suspension. After mixing the suspensions, the pH values of the oxalate suspensions were still at 7.60 f 0.01. Initially, the diluted suspensions each contained approximately 4.4 X M total Mn, primarily as suspended Mn304,and their oxalate concentrations were 0.00, 0.76, 1.25, 2.90, and 5.00 X M. The diluted suspensions were aged at 25.0 f 0.1 OC in a shaker bath while passing air over them and exposing them alternately to 12-h long-wave UV and 12-h darkness. The light intensity was comparable to that used in engineering practices with regard to exposure of building materials to sunlight. After 145 days, the shaking, flushing, and UV exposure were discontinued, and the diluted suspensions were left in the dark until the final analyses at age 1 year. Another set of solutions was prepared that contained M 0.01 M Na2S04,no Mn, and 0.65, 1.25, and 2.50 X oxalate and that had a pH of 7.60 f 0.01. These nonmanganese solutions were treated the same as the diluted suspensions and were used as a check for evaporation and oxygen oxidation of oxalate. Precautions Applied. The suspensions were prepared and maintained in an abiotic condition. All dissolved chemicals were passed through 0.10 Mm diameter pore Nuclepore filters, with the exception of NaOH, which was considered already abiotic. (Use of brand names in this paper is for identification only and does not constitute endorsement by the US. Geological Survey.) The water used was distilled, demineralized, and autoclaved. All Environ. Sci. Technol., Vol. 22, No. 1 , 1988
03
Table I. Oxidation Numbers of Manganese Oxides in the Oxalate Seriesa
initial oxalate concn, M 0.00 7.60 12.5 29.0 50.0
10
24
2.76 2.73 2.74 2.72
2.89 2.75 2.73 2.74
age, day 67 109 2.93 2.81 2.70 2.68 2.76
3.03 2.82 2.68 2.72 2.68
145
365
Results and Discussion
3.03 2.85 2.69 2.69 2.71
2.98 2.83
The formation of y-MnOOH via Mn304can be represented as a two-step cyclic, irreversible process. First, the Mn(I1) is oxidized to form Mn304and then Mn(II1) in the Mn304is protonated to form y-MnOOH along with the release of Mn(I1). The released Mn(I1) oxidizes to form more Mn30,. These two steps are diagramed in eq 1 and 2 (1). The oxidation rate is proportional to the surface Mn2++ 1/602(aq)+ H 2 0 1/3Mn304+ 2H+ (oxidation) (1)
2.72
“The oxidation number of Mn30, is 2.67 and that of y-MnOOH is 3.00. Intermediate values represent a mixture of these two oxides. At 3 days, all suspensions had an oxidation number of 2.67 f 0.05.
equipment in contact with the solutions were autoclaved or exposed to short-wave UV. Just before measuring the pH, the electrode was washed with 95% ethanol and dried. The flushing air was filtered through a sterile 0.45 pm diameter pore Gelman Acrodisc filter to remove air particulates, through a Gelman activated carbon capsule to remove organics, through a sterile 0.20 pm diameter pore Gelman pleated capsule filter to remove any carbon that might possibly have passed on from the carbon capsule, and through 0.01 M Na2S04to minimize evaporation of the suspension solutions. To insure that all samples were flushed at the same rate, the flasks were connected in series. The samples were arranged in order of increasing oxalate concentrations for the Mn-free solutions and then decreasing oxalate concentrations for the Mn suspensions. This arrangement minimized effects of possible passage of material to succeeding flasks. After extensive aging with continued shaking and flushing, solid material dried on the sides of the containers containing the suspension, and the airstream carried some of these particles over to the next vessel. Because of such carry-over, after 145 days a small amount of oxalate could be detected in the originally non-oxalate suspension. Evaporation was estimated from the sodium concentration increase. The amount of oxide in each suspension was estimated by subtracting the total dissolved Mn from the initial total Mn. In this paper the term “total dissolved Mn” represents Mn remaining in the filtrate after it was passed through a 0.1-pm filter. Effect of Light on Manganese Oxidation. The remaining undiluted Mn304 suspension was divided into equal portions and aged, unstirred, with one portion in the dark and the other exposed to indirect north daylight filtered through the laboratory window. These suspensions were in a 0.01 M Na2S04medium and at a pH of 7.82 when divided into portions to age. Analyses of the Test Solutions and Suspended Matter. Periodically, pH values and Mn and oxalate concentrations of the solutions were measured and oxidation numbers, X-ray diffraction patterns, and electron micrographs were obtained for the solids. The pH value was measured with a combination galss-calomel electrode and a Radiometer PHM 84 pH meter. The solids were separated from aliquota of the aging suspensions by passing the solutions through 0.10 pm diameter pore Nuclepore membrane filters. Oxidation numbers were determined by dissolving the solids remaining on the membrane filters in a known amount of oxalic acid solution to which sulfuric acid was added. Then an aliquot of each acidified, dissolved solids solution was back-titrated with standardized KMn04 (12).Mn in the untitrated portion of the dissolved solids solutions and Mn in the filtrate that had passed through the 0.1-pm filter were both determined by direct flame atomic absorption (AA) spectrophotometry. Ex64
Environ. Sci. Technol., Vol. 22, No. 1, 1988
perimental uncertainty in the oxidation number is f0.05. Oxalate in the filtrate and in the Mn-free solutions was also determined by KMnO, titration,
Mn304+ 2H+
-
-
2MnOOH
+ Mn2+ (protonation)
(2)
area of solid per unit volume (48). Therefore, the suspensions were prepared with the same initial amount of solid material in each flask and with a sufficient quantity of solid material so that changes in total surface area during the “cyclic process” would be negligible. Mineralogy of Suspended Solids. The oxidation numbers of the suspended manganese oxides in all the diluted suspensions at day 3 were 2.67 f 0.05, that of Mn304. By day 67 the oxalate-free suspension’s oxidation number was 2.93, close to that of y-MnOOH, whereas that of the 7.6 X M oxalate suspension was 2.81, halfway between that of Mn304and y-MnOOH, and those of the 1.25 X M and more highly concentrated oxalate suspensions remained close to that of Mn304. Even after 1 year of aging, the solids in the oxalate-containing suspensions had not increased in oxidation number, but that in the oxalate-free suspension had an oxidation number of 3.0. Table I shows the oxidation numbers of the suspended solids at the different ages tested. The 1-year-old precipitates from the non-oxalate suspension were long, well-defined needles characteristic of y-MnOOH, and those from the 7.6 X M oxalate suspension were a mixture of smaller needles and of the Mn304pseudocubic form. Those from the 2.90 X M oxalate suspension were mostly the pseudocubic form with a minor amount of stiU smaller needles. These precipitates are illustrated in Figure 1 (parts A-C, respectively). In regard to X-ray diffraction, established literature values of d spacings and intensities of reflections for yMnOOH (14) and for Mn304(49) compared well with those for the suspended solids. The pattern for the solids from the oxalate-free, 145-day-old suspension consisted of 13 reflections characteristic of y-MnOOH and only one characteristic of Mn301. Compared with the 100 intensity d spacing of y-MnOOH (3.41 A), the single Mn304 reflection (1.54 A) had an intensity of -1. The pattern for the solids from the 7.6 X M oxalate suspension consisted of 11 Mn304reflections and 10 y-MnOOH reflections, and the pattern for the solids from the 2.9 X lom3 M oxalate suspension consisted of 11 Mn304 and 4 yMnOOH reflections. For the solids from the 7.6 X M oxalate suspension, the intensity of the major y-MnOOH reflection (3.41 A) was 93% of that for the major Mn304 reflection (2.48 A), and for the solids from the 2.9 X lom3 M oxalate suspension, the intensity of the major yMnOOH reflection was only 8% of that for the major Mn304reflection. Thus the oxidation numbers, electron micrographs, and X-ray diffraction all suggest the same mineral identities and same relative amounts of each mineral. Activity Relations. Activities of the Mn and oxalate species were calculated by application of pH values; sulfate,
A
."-
3.0 3.5 4.0 4.5 5.0 5.5 6.0 -LOOT o l d Oxolola Cancentrotion ( n o i i ~ )
Flgure 2. Ratios 01 uncomplexed Mn(I1) and of total manganese oxalate [MnC20,' plus Mn(C,O,),2-] actinies to total Mn activity for different total oxalate end total Mn concentrations near neutral pH, at 0.01 imic sbecgth, and at 25 "C. Most of t b wmplexed Mn is in the MnC,O: state [fuexample. when the activity ratios of MnC,O:ltotal Mn are 0.917. 0.878. 0.737. and 0.467. the activity ratios 01 Mn(C20.)22-ltotalMn are 0.083.0.026, 0.007, and 0.0015. respectively]. log 01 total Mn Concentration: (0) -3.00; (0) -3 50: (A)-4.00 (0) -5.00 and -7.00. Ionic Strength = 0.01.
1. lyear-old Mn precipitates from the oxalate series. innial oxalate: (A) none: (B)7.6 X lo-' M; (C) 2.9 X W3M.
+re
Mn, and oxalate concentrations; and activity coefficients to equilibrium constant expressions. Equilibrium constants were calculated from free energies or taken from the literature (I3,50-52). The activity coefficients were calculated from the extended Debye-Hiickel law and ionic strengths. Because there was little or no evaporation through day 67, the ionic Strengths of the lo-, 24-, and 67-day-old Suspensions were assumed the same as those of the freshly prepared suspensions. The ionic strengths for the 1-year-old suspensions were calculated considering the sodium concentration increase to be a measure of Na2S04concentration increase. By 1year, the total Mn
(suspended plus dissolved) in the 1.25 X l0-l M oxalate suspension and the amount of evaporation from this flask differed considerably from the other suspensions, and there was no more 5 X lo4 M oxalate suspension. Thus, the 1.25 X 1O-l and 5.0 X 1O-l M oxalate suspensions were only considered through day 67. The HC20, activity is too low to be included in the activity calculations. The ratio of HCz04-activity to Cz04" activity is 0.01 at pH 6.27. This ratio decreases with increasing pH. When the diluted suspensions were prepared, the dissolved Mn species activities changed from preexisting values, and in the presence of oxalate, manganese oxalate complexes formed. As the Mns04 was protonated (eq 2), the Mn(I1) was released, and a portion of the released Mn(I1) was complexed by oxalate instead of directly reoxidizing to form more Mns04 (eq 1). The increased activity of the major manganese oxalate species MnCz02 became sufficient to more than compensate for the decrease in the dissolved, uncomplexed Mn(I1) and MnSO: activities. As a consequence, the total dissolved Mn activity increased in relation to the initial oxalate concentration. The activity values for the 10-day-old suspensions are listed in Table 11. As the suspensions aged, the total Mn activities decreased in all suspensions. By l year the oxalate-free suspension had only about one-third and the 2.9 x M oxalate suspension had a little over one-half the total Mn activity they had a t 10 days. Evidently, precipitation continued to occur in all suspensionsthroughout the aging period. Near neutral pH and a t 0.01 M ionic strength and 25 "C, in lo" M (0.55 ppm) and less Mn, the proportion of Mn activity complexed by oxalate varies only with oxalate concentration, but a t higher Mn concentration the proportions of Mn activity complexed varies also with the Mn concentration. Roughly 50% of Mn activity exists as manganese oxalate in concentrations of l(r M oxalate and of M and less Mn. Figure 2 illustrates the molar concentration ranges of total Mn and total oxalate over which the proportion of manganese oxalate activity is sufficient to noticeably contribute to the dissolved Mn activity at near neutral pH, at 25 "C and a t 0.01 M ionic strength. Reaction Affinities. Reaction affinities thermodynamically evaluate whether a reaction may occur. If the reaction affinity is negative, the process is impossible, but if the reaction affinity is positive, the reaction is feasible. Environ. Sci. Technol., Vol. 22, No. 1, 1988 65
Table 11. Activities and Reaction Affinities for the Oxalate Series at 10 Days XIO-5
initial oxalate concn
MnC204"
Mn(Cz04)z2+
activity MnS0,O 2.87
0.0
76.0 290 500
M
4.03 5.47 13.94
0.05 0.35 0.95
1.10
0.63 0.60
However, even if the reaction is feasible, kinetic factors may prevent the reaction from occurring a t a significant rate. The expression for reaction affinity a t 25.0 "C and 1 atm pressure is
A = -1.364(10g Q - log K ) (3) where Q is the activity quotient and K is the equilibrium constant. To determine whether or not the transformation from Mn304to y-MnOOH is favored by the observed decrease in Mn activity and pH, reaction affinities were calculated for Mn(I1) oxidation to Mn304and for Mn304 alteration to y-MnOOH. Reaction affinities were also calculated for MnCz04.2H20precipitation and for oxalate reduction of Mn304and of y-MnOOH. The reaction affinities for the suspensions at ages 10, 24, and 67 days and a t 1 year were calculated for the reactions described by eq 1and 2 and are designated Al and A2, respectively. At 10 days, AI decreases and A, increases with increasing initial oxalate concentration. These activity trends predict that the greater the concentration of oxalate, the more the oxidation reaction is retarded and the more the protonation reaction is encouraged. Initially, the protonation reaction may have been encouraged by the MnC204complex formation, and the concentration of this complex may have been sufficient a t the solid surfaces where these reactions occur to lower the favorability for the oxidation reaction. The presence of y-MnOOH crystaIs in all test suspensions showed that some Mn(II1) from the Mn,04 solid had been protonated (Figure l), and the increased total dissolved Mn(I1) in all oxalate suspensions indicated that when oxalate was present not all released Mn(I1) was immediately reoxidized. From day 10 through day 67 the reaction affinity trends persisted, and most of the reaction affinities did not change by more than 0.4 kJ mol-l. Regardless of the continuing prediction of greater favorability for conversion to yMnOOH in the presence of oxalate, the oxalate-free suspension formed y-MnOOH more rapidly than suspensions with oxalate present. By day 67, the solids from the oxalate-free suspension already consisted of -80% yMnOOH, and solids from the 7.6 X M oxalate suspension consisted of -43% y-MnOOH, while solids from M and higher oxalate suspensions continthe 1.25 X ued to contain the amount they had at day 10, -15% y-MnOOH. By day 109, the solids in the oxalate-free suspension were completely converted to y-MnOOH. The wide pH range at 1year (7.44-7.76) indicates that neither solution conditions nor calculated reaction affinities for these aged suspensions were comparable. Regardless, the average oxidation state, the electron micrographs, and the X-ray diffraction patterns indicate that the character of the 1-year-old suspended solids in the 7.6 X and 2.9 x M oxalate suspensions had not changed further. Evidently, oxalate somehow interfered with the reaction cycle. Figure 1illustrates blockage of the reaction cycle in the presence of oxalate, and Table I1 lists the reaction affinities a t 10 days. Besides the possible influence of a oxalate complex, the 66
Envlron. Sci. Technol., Vol. 22, No. 1, 1988
reaction affinity, kJ mol-' Mn2+
CMn
2.83
5.70 6.28 10.12 16.15
1.11
0.67 0.66
CZO4'-
pH
AI
A2
40.0 139.4 230.6
7.38 7.40 7.45 7.44
20.17 18.03 17.36 17.24
1.97 4.10 4.77 4.90
reducing ability of oxalate must also be considered. Even though reaction affinity values indicate that reductions of Mn304and y-MnOOH are thermodynamically feasible, other factors, probably kinetic in nature, prevented observable reduction through day 67. Oxalate reduces synthetic manganese(II1,IV) oxides under acidic conditions, and the rate of reduction decreases with increasing pH. Above pH 7.0, the reduction rate is too slow for flame atomic absorption spectrometric measurement of the total Mn released (11). In this experiment most of the oxalate remained in the solution through age 1 year, but after day 67 estimation of oxalate oxidation could not be made with confidence because oxalate concentrations were slightly altered by both evaporation and flask to flask transport. The constancy of the oxidation numbers and the total dissolved Mn(I1) decrease in the aging oxalate suspensions suggest that reduction via oxalate must have been minimal. Thus, action of a MnC2040complex at the surface or supersaturation with respect to MnCzO4-2Hz0remains as possible explanation of the observed data. Saturation with Respect to MnC204-2H20.Manganese oxalates with zero, one, and two water molecules have been characterized by X-ray diffraction (49). The precipitate expected in aqueous suspensions is the dihydrate. The solutions were all undersaturated with regard to this solid. There is probably a higher concentration of Mn(I1) at the solid surfaces compared to solution Mn, which is spatially separated by water and other solute molecules. Considering the possibility that saturation at the surface may have been approached, the solid surfaces were tested for a coating of MnC204.2H20. X-ray diffraction patterns showed no peaks relating to this material, and there was no analytically detectable oxalate attached to the solid surfaces. Either there was no such oxalate precipitated on the surfaces or the precipitate layer was too thin to detect. Figure 3 diagrams the solubility of MnCzO4.2H20 and the unsaturated condition of the diluted suspensions in terms of oxalate ion activity and total Mn activity. Possible Surface MnC204. At the pH used here, the average Mn304surface charge should be negative. This conclusion is based on the report that the zero point of charge for Mn304has been measured at pH 5.4 in 0.01 M NaCl and a t pH 5.7 in 0,001 M NaClO, (53). In addition to this average negative charge, there may be localized positive charges due to "surface manganese". The surface manganese could consist of Mn that is still part of the solid surface, adsorbed Mn, or Mn newly released by the protonation reaction. A possible manganese oxalate surface structure in the form of a hybrid of the manganese oxalate dihydrate and the dissolved manganese oxalate chelate may be postulated from consideration of some established features as analogies: (1)The greatest stabilization of a resonating structure is achieved when the contributing structures are energetically equivalent (54). For instance, in the case of the carboxyl group, the double bond may attach to either oxygen, but the most stable structure is one with reso-
-2.0 c1
plrj
i
PO
"surfac
-3.0
,.+Qaflo,
*
-
SurfaceMe
f
32' -4.0
0 A
-
A
0
0
c
c"
-.
f -5.0
4
L
L -2.0 -3.0 -4.0 -5.0 -Log C ~ 0 4 ~ ACTIVITY -
Figure 3. Total Mn activity for a solution saturated with MnC201.2H20 near neutral pH, at 0.01 Ionic strength, and at 25 'C and for the solutions from the oxalate serles at 10 days and at 1 year all relative to oxalate activity. The line is saturation with respect to MnC20,.2H,0. At age 10 days (0)= 5.00, (A)= 2.90, and (0) = 0.76X lo3 M hklal oxalate and at age 1 year (A)= 2.90, (0)= 0.76, and (0)= 0.00 X M initial oxalate.
,.,.%@., I
Fe
. 0@dC%/
,o%po%,
I
Fe c"&C\o"
+ .,o .. Fe
I
etc.
"-o**C*.O"
continuum. Water is coordinated to Fe from
Figure 4. a-FeC,O, above and below.
"I- [ q-
nating electrons equally distributed between the two oxygen atoms as shown:
[ p. -c
e.-
4.
-c
-c,
?d*
\o
Dotted lines represent resonating electrons in the illustration above and also in Figures 4 and 5. (2) A chelated ring structure is much more stable than a straight chain, single-bond structure (54). (3) Five- and six-membered rings are generally the most stable ring structures (54). (4) Solid a-FeC204.2H20,the more organized iron oxalate dihydrate, has been shown to be a continuum of (FeC204),(Fe is bonded to an oxalate on each side and is coordinated to water from above and below) (55, 56) (Figure 4 ) . ( 5 ) Because a-MnCz04-2H20is isostructural with aFeC204.2H20,Mn should be located in the same position as Fe (56). As opposed to the dihydrate solid, each Mn(I1) in the dissolved manganese oxalate complex has both bonds chelated by the same oxalate group and is surrounded by coordinated water molecules. The surface structure may be formed by the following steps: (1)As in the dihydrate solid, the dissolved dissociated oxalate ion may chelate Mn via both carboxyl groups to form a five-membered resonating ring. A MnC2040ion may become attracted to a charged surface and become polarized. This ion may originate in the bulk solution or may be freshly formed by the reaction of oxalate with Mn(I1) that is released when Mn(II1) is protonated (eq 2). The positive Mn(I1) end of the polarized ion can attach to the average negative Mn304surface charge. The surface manganese oxalate structure would have an unsatisfied negative charge on the exposed oxalate.
%@... 1 Me+ *,,
negatlve surface
\
--..&C\a,*
%&"..
I*
SurfaceMe
0 0
+
1
'...o.rC*.o"~
+
E
nearby harges
ive
anganese"
Me+f
"surfa
manganese"
I
...o*ICrpo,/ 0
Figure 5. Postulated mechanism for the bindlng of oxalate to surface manganese. The symbol "Me"may represent a divalent metal ion but in the present experiment represents Mn.
(2) The dissolved oxalate ion would be most stable as a resonating structure with a charge on both sides. If the Mn had only one charge exposed, such as Mn(I1) bound to the surface by one bond, one of the negative oxalate charges could attach to surface manganese and again the surface manganese-oxalate complex formed would have a residual negative charge. (3) A positively charged metal ion could attach to this unsatisfied negative charge. However, if this second metal ion were also attached to the solid surface, a less stable, bent resonating structure would be created. In contrast, a more stable, straight resonating structure would exist if the second metal were on the solution side of the oxalate unit. (4) If the dissolved Mn-oxalate chelate were polarized by a positive surface manganese, the negative end of the polarized unit would attach to the positive surface manganese. The final postulated surface manganese-oxalate-Mn structure, regardless of formation mechanism, would be a hybrid between the dissolved complex and the Mn-oxalate dihydrate solid. Figure 5 is a schematic of these reaction mechanisms. Steric hindrance of the Mn(II)-Mn304-yMnOOH cycle would occur as the solid surfaces become covered with oxalate structures such as A and B in Figure 5. In Figure 5, Me stands for Mn but Me may represent other divalent metals such as those in the isostructural series of a-Me oxalate dihydrates (Me = Mg, Mn, Fe, Co, Ni, and Zn) (56). These dihydrates have been structurally analyzed (57),have been reported as minerals, and are likely to occur where oxalic acid excreting lichens colonize suitable substrates (34). The metals in these dihydrates are common in natural waters and in the materials these waters contact and would compete with Mn in this hybrid layer. Auger electron spectrometry perhaps could indicate whether there is carbon on the particle surfaces and whether there is an abrupt delineation of decreased carbon concentration below the particle surfaces. Comparison of precipitates containing primarily Mn304with those containing a much greater percentage of y-MnOOH might indicate whether the carbon is only on the Mn304or on both the Mn,04 and the y-MnOOH precipitates. Also, perhaps mass spectrometry might detect the presence of oxalate associated with the precipitates, and electrophoresis may determine the polarity of the surface charge on the particles. Other carboxylic acids may attach to surface manganese. The strength of the attachment and whether a redox reaction occurs depend upon several factors, i.e., the charEnviron. Sci. Technol., Vol. 22, No. 1, 1988 67
acter, location, and number of functional groups in the organic molecule, and the resonance of electrons within and between these groups, especiallythe -COOH group(& Other functional groups may be -OH, -OCH, -CH,, -NO3, -C1, -Br, -I, etc. Compared to other positions on the organic molecule, if these groups are attached to the same carbon atom as the most acidic -COOH group or if located in the ortho position of an aromatic carboxylic acid, they will generally have a greater influence on the chemistry of that -COOH. The influence depends on the capacity of the functional groups to attract and to repel electrons. The carbon to which the -COOH groups are attached is often part of the electron flow path in resonating hybrids of organic compounds. The resonating atoms of the hybrids are the focal point for many reactions, including chelation and redox reactions (54). Natural waters and effluents frequently contain organic materials with some combination of these features (58). The reactive portion may involve all or part of a specific organic compound or be a portion of a fulvic or humic acid molecule. If there is more than one locale of negative charge of sufficient strength and if there is do steric hindrance, a Me(I1) also may attach to the solution side of the coating. As stated above, the proximity of other functionalgroups determines the reactivity and consequently the Me binding capacity of a -COOH group. The carboxyl groups in aliphatic dicarboxylic acids act independently of each other when separated by five or more carbon atoms and have acidities similar to acetic acid. As the number of carbon atoms becomes fewer, the inductive effect of one carboxyl group on the other causes the acid strength of the first dissociated carboxyl group to be greater than that of acetic acid and the acid strength of the second to be less than that of acetic acid. A different character for oxalic acid is caused by the lack of carbon atoms between the carboxyl groups, permitting close proximity of the resonating carboxyl groups (54). In aliphatic dicarboxylic acid crystals with the composition (CH,),(COOH),, where x = 1-6, the molecules are bound together end to end via hydrogen bonding between the carboxyl groups. This is true even for certain isomers with one or two -CH3 groups or a -C=O group substituted for hydrogen(s) on -CH2- unit(s) (59). This end to end bonding may also exist for larger molecules. Mn(I1) in such a dicarboxylic acid Mn(I1) precipitate would probably take the place of the two hydrogens and bond to carboxyl groups from two different molecules, forming two adjacent four-membered resonating rings. Again the contribution of the other functional groups in the acid partially determine (a) whether the acid is capable of binding with a surface manganese, (b) the number of members in the surface manganeseorganic acid resonating ring (and consequently the stability), and (c) whether a solution Mn(1I) can attach to the solution end of the attached acid. Influence of Turbulence and Light Exposure. That turbulence of a water system can influence crystal formation and ultimately particulate transport is demonstrated by the differences between shaken and undisturbed suspensions containing no oxalate. By 145 days, the undisturbed suspensions had grown long, fine, hairlike needles that sould easily become entangled with each other or any barrier in the flow path of slowly moving water, whereas, the shaken oxalate-free suspension had grown much shorter y-MnOOH needles by this age. The 1:1 dilution may have slowed needle-length growth, but more likely the vigorous shaking broke up the needles or prevented localized high surface concentration effect by 68
Environ. Sci. Technol., Vol. 22, No. 1, 1988
keeping the solids dispersed. Other workers have shown light to affect the oxidation-reduction reactions between manganese oxides and organic matter. In the two concentrated, undisturbed, nonorganic suspensions, the presence of indirect sunlight filtered through window glass as opposed to complete darkness made no measurable difference in the oxidation pathway from day 3 when the oxidation numbers were 2.81 f 0.05 to day 145 when the oxidation numbers were 3.04 f 0.05. (The persistence of y-MnOOH is shown by the oxidation numbers that were already 2.95 by day 24, 3.02 by day 67, and still 3.04 at day 145.) Figure 6 demonstrates the needle-length differences between the shaken and unshaken suspensions and the similarity of the suspensions aged, one in darkness and one in subdued daylight. In the case of the oxalate series, even though precipitation continued through age 1year, the quiet aging in the dark for the last 0.5 year did not alter the mineralogies of the precipitating solids from that established earlier when .exposed to agitation, aeration, and alternately to 12-h long-wave UV and 12-h darkness. In other words, the influence, if any, on the reaction mechanisms or reaction rates due to UV occurs during the early stages of the reactions described here.
Conclusions A natural organic compound, oxalic acid, capable of reducing Mn at low pH values and possessing little or no such capability at pH values just above neutrality may still alter the reaction pathway and change the character of the manganese oxide surface. For a period of 1year, oxalate retarded the “Mn304to y-MnOOH cycle” in a 7.6 X M oxalate suspension and completely blocked this cycle in 1.25 M oxalate and higher suspensions. (These suspensions were well aerated, shaken in a 0.01 M Na2S04 medium, and maintained at pH 7.4 f 0.2 and 25.0 f 0.1 “C.) During the first 67 days, neither oxygen nor the suspended Mn304and y-MnOOH oxidized oxalate nor did oxalate reduce these manganese oxides to any degree. Oxalate complexed so much free Mn(I1) that the total dissolved Mn(I1) concentration increased by 2 or 3 times in the M oxalate suspensions. This increased dissolved Mn(I1) concentration may aid in dissolved Mn transport and may act as a reservoir for reactions that require available dissolved Mn(I1). Whether this reservoir is biologically advantageous depends on nutritional requirements, toxicity sensitivies, and the capacity of the Mn complex to pass through the cell wall. The oxalate complexing of free Mn(I1) initially may have encouraged y-MnOOH formation. However, the oxidation numbers, electron micrographs, and X-ray diffraction patterns of the suspended solids indicated that, after aging, oxalate retarded y-MnOOH formation even though reaction affinity values continued to predict that oxalate could encourage y-MnOOH formation providing there was nothing to hinder the process. There was no experimental evidence to indicate that a precipitate of MnCzO4-Hz0had formed on the particulates. The solutions were unsaturated with regard to this material. Some form of manganese oxalate complex a t the solid surface may sterically block the transformation process. A suggested possible mechanism of coating formation illustrates how an organic material may alter mineral surfaces and create new sites to which metals may attach. Many naturally and anthropogenically produced organics have -COOH groups that may chelate surface manganese. In the presence of these organics, if solution chemistry and kinetics of the competing reactions with other dissolved species permit and if there is no steric hindrance, an or-
of the particle surface would interfere with the Mn@y-MnOOH cycle. Many divalent metal ions common to natural waters and to material that these waters contact may compete with Mn for sites on the chelated coating. The coating could consist of surface Mn(I0-organic material-solution Me(I0-charged hydration shell. Because, in natural waters, redox conditions and light intensities vary with locale and especially with depth, settling rates and transportability of the manganese oxides may determine further oxide alteration. If slowly moving waters contain no organic matter with properties similar to that of oxalate and have favorable chemical and physical solution parameters, long, hairlike y-MnOOH needles may form. Once the alteration of y-MnOOH is under way, the reaction process will not he changed, even if subdued daylight becomes no longer availahle. These long needles will more readdy become entangled with each other or with other surfaces that they may encounter. They may settle out and if so may encounter sufficient Eh-pH changes to become reduced to soluble Mn(I1). In contrast, if these slowly moving waters do contain a sufficient concentration of oxalate-type organic matter, finely divided Mn304will he the form of Mn precipitate that persists. If the organic matter concentration is decreased, y-MnOOH formation may resume. However, if the waters also become turhulent, the finely divided, unconverted Mn,04 and the newly forming, short y-MnOOH needles would be more easily kept suspended and transported than the elongated, hairlike y-MnOOH needles formed in quiet waters. In the presence of organic matter with properties similar to oxalate, once the ratio of Mn304-y-Mn00H becomes stabilized, if then deprived of long-wave W light, these oxides would continue to precipitate with the established oxide ratio until other solution parameters change. Acknowledgments This work was done as part of the research program in aqueous chemistry of the Water Resources Division, US. Geological Survey. Helpful comments were furnished by my Geological Survey colleagues C. E. Roberson and J. A. Davis. Registry No. Hausmannite, 1309-55-3; manganite, 1310-981; oxalic acid, 144-62-7. Literature Cited
(3) Khan, T. R.; Langford,C. H.; Skippen, G. B. Org. Geochem. 1984, 7(3/4), 261-266. (4) Wangersky, P. J. Mar. Chem. 1986, 18, 269-297. (5) Martin, J. M.; Meybeck, M. Mar. Chem. 1979,7,173-2%. Figure 6. y-MnOOH needles at 145 days: (A) aged continuously shaken and exposed alternately to 12-h long-wave UV and 12-h darkness. but (B) and (C) both aged unshaken. (E) exposed to indirect daylight. and (C)exposed to complete darkness.
ganic coating on the Mn30, precipitate may form. Sufficient concentrations of these organic materials are most likely in locales of decomposing plant debris, of high hiological activity, or where certain effluent wastes are being discharged. Resonance of the chelated portion and of other functional groups in the organic coating may expose reactive sites with a different character than that of the original mineral surface. Mn(I1) from solution may attach to these new sites. Regardless of whether or not solution Mn(I1) was taken up by the new surface, the altered nature
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Received for review December 23, 1986. Accepted August 12, 1987.
