Oxidation of manganese by spores of a marine bacillus: Kinetic and thermodynamic considerations

Geochimica a Cosmochimica Ado Vol. 50, pp. 1819-1824 D Pergamon Journals Ltd. 1986. Printed in U.S.A.

Wl6-7037/86/53.00

+ .OLl

Oxidation of manganese by spores of a marine bacillus: Kinetic and thermodynamic considerations* DAVID HASTINGSand STEVENEMERSON School of Oceanography, WB-10, University of Washington, Seattle, WA 98195, U.S.A. (Received November 18, 1985; accepted in revisedform May 22, 1986)

Abstract-The catalytic properties of spores of a marine Bacillus known to oxidize divalent manganese were used to perform laboratory Mn(I1) oxidation experiments at environmental conditions of pH and Mn(II) concentration. We found that at pH 7.8 the initial kinetics of Mn(I1) oxidation facilitated by the spores was four orders of magnitude greater than that which would be expected for abiotic autocatalysis on a colloidal MnOz surface. The rate progressively decreased as the spores became coated with manganese oxide, eventually becoming very near that predicted for abiotic surface catalysis. Transmission electron microscopic observations and oxidation state measurements of solids precipitated at pH 7.5 and [Mn(II)] < 50 nM indicated that the initial oxidation product was hausmannite (Mn,04 or MnO, where x = 1.33) which aged to more highly oxidized Mn02 (x = 1.9)in the time scale of weeks. By utilizing spores to catalyze the oxidation rate, we were able to maintain our experimental system within the seawater range of pH and Mn(II) where highly oxidized manganese oxide precipitates are thermodynamically stable. In doing so we obtained, for the first time, laboratory precipitates with oxidation states similar to that found in marine particulate material. These results suggest that the concentration of manganese in seawater and the oxidation state of marine manganese oxidesare controlled by the rapid precipitation of Mn904, which can be microbially mediated, followed by the disproportionation to MnOz

Mn304 + 2H+ ++ 2yMnOOH

INTRODUCTION

THE GEOCHEMISTRY OF manganese in the marine environment is greatly influenced by redox transformations in the oxygen minimum zone (MARTIN and KNAUER, 1984), at hydrothermal vents (WEISS, 1977) in sediments (KLINKHAMMER, 1980), and in anoxic basins and fjords (EMERSONet al., 1979). Field studies in anoxic basins have shown that the oxidation residence time of dissolved Mn is much shorter than expected if it occurred solely by an inorganic mechanism (EMERSONet al., 1979). This evidence led to the suggestion that Mn(II) oxidation may be bacterially catalyzed. Subsequent field studies confirmed this microbial activity in several suboxic environments (EMERSON et al., 1982; TEBO et al., 1984; TEBO and EMERSON, 1985). The oxidation reaction of Mn(II) to MnOz can be summarized as: Mn+* + R02(aq) + Hz0 t* Mn02 + 2H+.

(1)

The mechanism is believed to be a two step process where Mn(I1) first oxidizes to a metastable intermediate, which should subsequently alter to a more stable, more highly oxidized solid (HEM, 1980; HEM and LIND, 1983). At 25°C the initial oxidation product of Mn(I1) is hausmannite, Mn304 : 3Mn+* + 3H20 + ~/zO~(~)c* Mn304 + 6H+.

(2)

Transmission electron microscopy, X-ray and electron difhaction indicate that Mn,04 spontaneously protonates to manganite, yMnOOH (MURRAY et al., 1985).

* University of Washington, Contribution No. 1639.

School of Oceanography

+ Mn+*.

(3)

In these experiments, which lasted 4 to 9 months, Mn304 aged to an oxidation state no higher than 3.0. Theoretically, Mn304 should disproportionate to form Mn02 : Mn304 + 4H+ c* Mn02 + 2Mnf2 + 2HzO.

(4)

However, reaction 4 has not been experimentally demonstrated. It is the established opinion that manganese oxides with an oxidation state above three do not form readily in laboratory systems except under extreme experimental conditions. By contrast, manganese in the ocean is typically highly oxidized; deep ocean Mn nodules are 98% Mn(IV) (PIPER~~al., 1984; MURRAY et al., 1984) the average oxidation state of Mn in marine sediments is 3.8-4.0 (MURRAY et al., 1984; KALHORNand EMERSON,1984) and particulate Mn above the 02/H2S interface in Saanich Inlet has an oxidation state as high as 3.6 (KALHORN and EMERSON,1984). HEM and LIND (1983) use the mechanistic arguments for Mn oxidation combined with free energy data to establish a non-equilibrium thermodynamic scheme that is described in Fig. 1. Since Mn02 is the most stable phase over most of the Mn(I1) concentration and pH range of the figure, this is not a true stability diagram; instead each reaction indicates a relative stability between phases. An equilibrium line is drawn for each reaction based on the appropriate equilibrium constants calculated from activities at standard temperature and pressure (02 saturation is assumed in reactions 1 and 2 and thermodynamic data are compiled from HEM and LIND, 1983). Based on these lines, and with Mn and pH values for a given solution, it is possible to predict which Mn phase is thermodynamically

1819

1820

D. Hastings and S. Emerson

HG. 1. Dissolved &In+*activity-pH equilibrium relationships in aerated solutions. Each line represents the equilibrium for the reaction written above each liae. Bared on these line5 and with Mn and pH vaIues it is possible to predict which Mn phase is tbezmo@namkUy

stable. A soiution with Mn

and pH values to the right of line 2 would be expected to p&&ate Mo,O+ A solution to the right of lines 3 or 4 would lie in the fiekl where Mn,O, is stable with respect to -rMnOOH or Mn02, rmpectively. Data points from experimen*rA,Bl,aodB2fromthissudy(~)aswtUasthrrseof HEMand LJND(1983; X) and MURRAYer ol. ( 1985; 0) are ShOWn.

stable. A solution with Mn and pH values to the right of line 2 is in the field where the energy for formation of Mn,O, has a negative free energy; from purely thermodynamic considerations the reaction should proceed to the right. For reactions 3 and 4 a solution to the right of the lines would lie in the field where Mn,O, is stable with respect to yMnOOH and where MnsO, is stable with respect to MnOz , nxpcctively. Data points from the previous studies of HEM and LMD (1983) and MURRAY et al. ( 1984), in which oxidation state values of 3.0 were attained, are plotted in Fig 1. Assuming the two step mechanism for Mn oxidation one would predict an initial oxidation product of MusO, (reaction 2) which ages to rhInOOH (reaction 3). Since all of these q.xpmimcntai points lie to the right of line 4, production of MnOz is not thermodynamically favored and would not be expected on thesegrounds.Asprrdicteatheobeavedsdidinthese expuiments was MnJOl which then altered to rMnOOH with no detecmble Mn02. The Mn concentration and pH field for seawater lies just above line 2 and below line 4 of Fig. 1 in the region where one would predict the initial oxidation product, MnsO,, to alter to the more highly oxidized MnO*. It could be, then, that the reason for the oxi-

dation state discrepancy between laboratory studies and field measurements is that the laboratory experiments have not been conducted at the appropriate pH and Mn(II) concentration. This is not simply due to an oversight on the part of the experimenters. but rather a constraint resulting from the very slow inorganic oxidation rate of Mn(I1); in order to derive a precipitate on an observable time scale either the Mn(l1) concentration or pH must be higher than in the natural setting The objective of this study was to form manganese oxides in a stability field near that of seawater in order to test the mechanism for manganese oxide formation and determine the oxidation product under environmentally realistic conditions. To achieve rapid oxidation kinetics at relatively low pH and manganese concentration we utilized the catalytic properties of spores of the marine bacillus SG- 1. SC- 1 is a salt-requiring bacterium isolated from a near-shore enrichment culture (NEALSON and FORD, 1980). The spores of this culture have been shown to bind and oxidize manganese extracellularly at pH values below which Mn(I1) amoxidizes; this oxidation is not coupled to energy metabolism since spores rendered non-viable still oxidize Mn (ROSSONand NEALEON, 1982). The catalysis is believed to be carried out by a spore coat protein that complexes Mn(II), which subsequently is rapidly oxidized. These spores provide an ideal microbial catalyst for studying the products of manganese oxidation, since they are highly resistant to physical and chemical changes and because they are dormant, which eliminates the need to sustain growth. We describe two experiments conducted in the pH and Mn(I1) range indicated in Fig. 1. Experiment A was a batch study designed to investigate the kinetics of manganese oxidation catalyzed by SG- 1 spores. This investigation is an extension of the work of ROSSON and NEAUON (1982) in which we include analysis of the oxidation state and mineralogy of the precipitate. The results are interpreted in the context of previously observed oxidation rate measurements in the laboratory and the environment. Experiment B was a constant composition study designed to observe the product of oxidation at a non-varying Mn concentration similar to that of the marine environment. METHODS Compbztelysponrlated cultof S-1 were harvestedand cleanedtookinapunsuqakon ofmature, dormant SpoKs accordiw to the method described by RCSON and NEALSON ( t982). The numba ofsporeswas detamioed by direct counts

under phase optics wiih a FetrofKHausaercounting chamber. To minimize span getmination the experiments were carried out in 0.2 firn 5hued seawater whieb was irradiated with UV light for 4 hours to remove d&lved organic carbon. Spores were phaar bright over the duratkm of the experiments indicating that t&y remained dormant and &d not germinate. In expetiment A two 11 bottles of seawater were stirred constantly and maintained at pH 7.8 and room tempuature (23°C). One bottle was kept anoxic by sparging it with a 0.1% CO2 in N2 mixturr; the other bottle was equilibrated with room air. Each bottle was inoculated with 3.2 X lO’/ml SCI spansfollowodby a seriesof d&bed Mn spikes es M&Q:

1821

Marine bacterial Mn oxidation 2 pM Mn(I1) as MnS04 was added on 4/17, 4/25, 6/19, 10 pM on 7123 and a final 2 PM spike on 1l/14. Dissolved Mn and particulate Mn bound by spores were determined periodically throughout the experiment by flameless atomic absorption spectrophotometry on a Perkin Elmer 5000 spectrophotometer with an HCfA 2000 graphite furnace and Zeeman background correction, or colorometrically by the formaldoxime method (BREWERand SPENCER,197 1). Experiment B consisted of two 2 1 bottles of 1.6 X lO’/ml spore suspension that were maintained at pH 6.5 and pH 7.6, respectively,by spargingwith CO2in sir mixtures. Each bottle received a constant input of 0.2 mM dissolved Mn as M&l2 approximately 2 ml d-‘, or about 200 nM d-‘. Mn input was discontinued after 15 hours in the pH 6.5 bottle and 25 days later in the pH 7.6 bottle. Dissolved Mn and pH were monitored periodically in each experiment. Particulate Mn oxides were characterized by transmission electron microscopy, scanning electron microscopy with elemental analysis, X-ray diffraction and oxidation state measurements. Crystal structure and morphology were determined by transmission electron microscopy on a JEOL 100-B electron microscope. TEM samples were prepared by washing three times in distilled HZ0 to remove salt, sonicating for l-2 seconds to disperse the agglomerated solids, and drying one drop of this Mn oxide suspension on a formvar coated carbon stabilized grid. Elemental composition of the Mn oxides was determined on a Cambridge MkIIA scanning electron microscope with a EDAX 707A analyzer. X-ray analysis was performed by the Debye-Scherrer method using a 114.6 mm diameter powder camera with Mn oxide filtered Fe K, radiation or V filtered Cr K,, radiation and 23 hour exposure times. Equivalents of oxidized Mn were determined by an iodiometric titration similar to a Winkler dissolved oxygen titration procedure adapted to especially low Mn oxide levels by reducing sample and reagent volumes by two orders of magnitude (MURRAYet al., 1984; KALHORN and EMERSON,1984). Ten to 60 ml of sample suspension were filtered through 13 mm 0.2 pm Nuclepore filter followed by 2 ml distilled Hz0 to rinse unbound Mn(I1) off the filter. Average oxidation state values are calculated from the number of oxidized equivalents and total Mn: average oxidation state =

oxidized equivalents + 2. total moles of Mn

Glassware and hlter holders were soaked in hot 1 N HCl for 6 hours, rinsed in distilled, deionized water, and dried in a laminar flow hood to prevent Mn contamination. The precision, determined by measurement of replicate samples, was kO.05 (10). RESULTS

AND DISCUSSION

No detectable Mn+’ was removed from solution to the spore surface in the anoxic suspension of experiment A; initial levels of dissolved Mn remained constant for two months. Rapid removal of Mn did occur in the presence of oxygen at pH 7.8 (Fig. 2). The data from these experiments is presented in the Appendix. The removal of Mn following the first two 2 PM Mn additions was complete after 48 hours. The removal rate after the third addition was significantly slower than in the first two. After the third spike, 10 PM Mn was added to coat the spore surface with Mn oxide. Following this, the removal rate for the final 2 PM Mn addition was substantially slower; 75% of the dissolved Mn remained in solution after 48 hours. The removal rate expression for Mn can be written as the summation of a biologically catalyzed removal

u n.zC

4

w 2

FIG. 2. (A) Removal of dissolved Mnf2 vs. time in experiment A. 2 rM Mn(I1) was added on 4/17 (spike 1, A), 4/25 (spike 2, 0), 6/19 (spike 3, A), 10 pM on 7/23 (spike 4; not shown) and a final 2 PM spike on 11/14 (spike 5, 0). (B) Average oxidation state vs. time for particulate Mn in spike 3, experiment A. The data from experiment A are tabulated in the Appendix.

with an inorganic rate expression similar to that for Mn oxidation by colloidal Mn02 (MORGAN and STUMM, 1971): dMn - = kb [spore sites] at +

k~,[Mn,l[OH-12tO~l[Mn’21(5)

where Mn, = particulate manganese and kb and kMo are the rate constants for biologically catalyzed and abiotic, autocatalytic Mn removal, respectively. For fresh spores, the biologically catalyzed rate is much faster than the inorganic removal rate. To compare the removal kinetics determined in this experiment with that of an inorganic system, a rate constant for each Mn addition was determined using only the rate expression for Mn oxidation by colloidal Mn02. Constant values of [OH-] = 10-6.2, [Or] = 230 PM with experimentally

determined

values for F,

[Mn+‘],

and particulate Mn result in the rate constants, k,,,, listed in Table 1. The rate constant for the initial Mn spike, ke = 1.6 X 1O22Mm4d-l, was almost four orders of magnitude greater than the rate of 5.5 X 10” which would be expected for the abiological MnOr system (BREWER, 1975). It approached the value determined

by EMERSONet al. ( 1979) in Saanich Inlet, where bacterial catalysis of Mn oxidation has been observed. As more dissolved Mn was added and subsequently removed from solution by the spores, the spore surface became increasingly more encrusted with Mn oxide and kMn decreased. Eventually in spike #5 the rate constant was almost identical to that predicted for an abiotic suspension of colloidal Mnb2: km = 5.1 X 10L8.

These data suggest that during spike #5 the spores were

D. Hastings and S. Emerson

1822 Table

spike

1

1.

Mn removal rate constant, kMn, calculated for spikes #l-5 in experiment 1. Average oxidation states of particulate Hn Immediately after each Mn addition are also shown.

date

4/11

HIl(II) added

2 UM

: 6119 4/25 2 UH !JM 4 7123 10 “M 5 11/14 2 UM removal on colloidal MnO, (Brewer, 1975) in situ, Saanich Inlet 7ziGrson ” g., 1979)

kfl” mol-‘L-‘d-l 1.6

x 10”

2.8 2.1

x 10” 10’0 _ ._ 5.1 x 10” 5.5 x 10’8 8

x 10’2

initial average oxidation state

2.8 3.3 3.4 3.5 3.5 --_--

completely covered with, or surrounded by, a Mn precipitate blocking the biologically induced catalysis. At pH 7.8 and a manganese concentration of 2 PM, the reaction mechanism and stability scheme of Fig. 1 would predict an initial reaction product of Mn304 that alters to -yMnOOH according to reaction 3. The difference between our experiment and those done previously is that manganese was rapidly removed from solution after each spike causing the Mn(II) concentration to drop into the field where Mn@, should disproportionate to MnQ and Mn(II). This sequence should result in a product with an oxidation state that is initially in the range of 2.7 (Mn304) and alters to a higher value in the range of 3.0 to 4.0 (yMnOOH and MnOz , respectively). The initial average oxidation states of the solid formed by the spores (Table 1 and the Appendix) indicate a progressive range from 2.8 to 3.5. Because these measurements are on mixed phases, unequivocal mineralogic information is ob-

sawed; the salient result is that an oxidation state of

3.5 was achieved near the end of the experiment. Immediately following each Mn addition there was a consistent trend for the oxidation state to decrease sharply then gradually increase. An example of this trend for spike three is iliustrated in Fig. 2b. The initial apparent negative oxidation rate reflects the rapid adsorption of Mn(I1) to the spore surface, which is followed by the slower oxidation of the adsorbed Mn. This interpretation is consistent with the initial rapid (O-l hr) removal of dissolved Mn after each of the Mn additions (Fii. 2a). The data suggest there are a limited number of adsorption sites, and once these sites are saturated, Mn(I1) removal rate deereases. An adsorption site is subsequently regenerated by the oxidation of the adsorbed Mn(II). Since initial adsorption is fast, whereas complete removal takes over two days, these data indicate that oxidation is the rate limiting step in the removal of Mn. Analysis of the Mn oxides by transmission electron microscopy following the final Mn(I1) spike indicated that there were very few crystalline structures resembling Mn oxides that have been previously identified. Only one group of Mn#& crystals was identified from a study of over two hundred fields at 17,000X magnification (Fig. 3a). No precipitates resembling the yMnOOHfoundbyHElrrandLIND(1983)orMunRA~ et al. (1985) were observed in this study. Much of the Mn oxide consisted of microcrystalline solids which were X-ray amorphous and not easily identified by electron microscopy. The most prevalent form was a crumpled, sheety material shown in Fig. 3b. Elemental analysis is SEM/EDAX showed that Mn was the only

Fro 3. Transmission electron micrographs of particulate h4n oxides. The characteristic morphology for Mn,O* is small octahedra. At the end of experiment A, one cluster of Mn@, crystals was identitied (a); the most prevalent form with chamcteristic features was a crumpled, she&y Mn oxide (b). At the onset of experiment B, Mn& was ubiquitous and was the only crystalline form observed (c). Bars represent 1Nrn.

1823

Marine bacterial Mn oxidation element over atomic number 9 to be detected on this material; it has characteristic features very similar to a Mn precipitate resembling vemadite or 6-Mn02 (TIPPING et al., 1984). The large range ofmanganese concentrations in experiment A causes some equivocation with regard to a thermodynamic interpretation of the results using the oxidation state measurements. To circumvent these difficulties, we performed two constant composition experiments in the pH and Mn(I1) concentration range indicated in Fig. 1 (experiments Bl and B2). The Mn(II) and pH ranges were chosen: (1) to test whether or not spores facilitate the direct precipitation of MnOa without the MnzO intermediate (experiment Bl, left of line 2 in Fig. I), and (2) to measure the oxidation state achieved in an experiment that is constrained to the MnOa stability field (experiment B2, between lines 2 and 4 of Fig. 1). Dissolved Mn accumulated rapidly in the reaction vessel of experiment Bl at pH 6.5; after 15 hours the Mn(I1) concentration reached 130 nM at which time the supply of MnOz was discontinued. No removal onto the spore surface was observed after 5 weeks (Table 2). Dissolved Mn levels decreased by approximately 10 nM during this period probably due to removal onto the glass walls of the flasks; total Mn measurements always equalled dissolved Mn measurements indicating that there was no measurable suspended particulate Mn. This experiment indicated clearly that at pH 6.5 the catalytic properties of the spores do not facilitate the direct oxidation of Mn(I1) to MnOz. In experiment B2 at pH 7.6, rapid Mn(I1) removal occurred over the 25 day period during which Mn(I1) was being added. While 200 nM Mn(I1) d-i was added, dissolved Mn levels remained below 30 nM after the first day. After 12 days average oxidation state values were over 3.5 indicating that highly oxidized (Table 2) Mn oxides were formed by the disproportionation of Mn304 to MnOz . After 25 days, when manganese addition was terminated, the average oxidation state remained at or near 3.8. Analysis of the initial precipitate by TEM revealed that Mn304 was the only crystalline form (Fig. 3~). As the solution aged, the relative proportion of Mn304 decreased and uncharactetized

Table 2. CumulativeMrl(II)added.cnn(II)l oxidation state and 82 (pH 7.6)

day

total

0 1

Mn(Il) Bl

0.0 0.13a 0.13

2 4 6 9 12 17

input

was

0.0 0.2 0.4 0.8 1.2 1.8 2.2

discontinued

for

and average

experiments

CMn(II)I, uM 81 82 .06

.1* _-

.05 .04 .Ol .Ol .Ol .02 .02 .02 .Ol __

.I2

.Ol

.I3 .10 .08 .I4 _.ll __

3.1 4.0b 4.1 4.1

23 25 26 'Mn

added,vM 82

values

after

15 hours

bM” input wasdiScOnti”“ed after 25

days

(pH6.5)

Bl

oxidation 82

__ __ __ 3.10 3.39 3.57 __ 3.70 3.70 3.85 in expt. in expt.

81. B2.

state

amorphous Mn oxides predominated. We observed no X-ray diffraction lines from any of these samples. The continuous addition of low levels of dissolved Mn(I1) at pH 7.6 yielded highly oxidized particulate Mn with at least 80% Mn(IV) or MnOz . By catalyzing the initial oxidation step, which facilitates the rapid removal of dissolved Mn from solution, we enabled the oxidation reaction to take place in a chemical environment where MnOl formation is thermodynamically favored. This region includes the range of Mn concentrations and pH values typical of fresh water and marine environments where Mn oxides are highly oxidized. It should be noted that the conditions of temperature and pressure used in these experiments are different from the deep ocean sea floor where many Mn(IV) oxides are found (BURNSand BURNS, 1979). CONCLUSIONS We have demonstrated the utility of using a microbial catalyst in laboratory experiments to derive unique information about the thermodynamic and kinetic properties of environmental manganese oxidation. While sporeformers are generally not considered to be the dominant bacterial flora in the marine environment, the presence of SG- 1 spores provided the catalyst necessary to perform laboratory oxidation experiments under environmental conditions on an observable time scale. Our experiments confirm the two step Mn(I1) oxidation mechanism whereby MnaO, is the initial reaction product, and they follow the thermodynamic prediction that this partially oxidized mineral should transform to a more highly oxidized manganese oxide at conditions similar to those of many natural environments. Furthermore, this reaction occurs on the time scale of weeks. The oxidation states of manganese oxides formed in the laboratory are very similar to those found in the aquatic environment if environmental conditions are maintained in the experiments. At oceanic pH values ranging from 7.8 to 8.1 the equilibrium concentration of Mn based on the Mn(II)Mn304 redox couple is predicted to vary from 0.3 to 2 nmol kg-’ (Fig. 1). Dissolved Mn concentrations in the open ocean typically range from 0.2 to 3 nmol kg-’ (BRULAND, 1983). This supports the mechanistic argument that Mn in the ocean is in equilibrium with Mn304 (KLINKHAMMERand BENDER, 1980). We suggest that Mn oxidation kinetics do not limit Mn(II)MnOa equilibrium in the ocean, as has been often suggested. Instead, initial rapid oxidation to a metastable intermediate, MnaO., , perhaps facilitated by microbial activity, controls the dissolved Mn concentration. Mn304 subsequently d&proportionates to the highly oxidized Mn02 typically found in the marine environment. Acknowledgements-We would like to express our sincere thanks to Dr. Bradley Tebo who taught us about spores and helped prepare the initial batch. We had numerous enlightening discussions with Dr. Tebo and he provided helpful sug-

1824

D. Hastings and S. Emerson

gestions throughout this research. We also appreciate the willingness of Dr. James Murray to share his knowledge about this subject and his advice as a critical reader of the manuscript. This research was supported by NSF Grant GCE 850277 1. Editorial handling: R. G. Bums REFERENCES BREWERP. G. (1975) Minor elements in seawater. In Chemical Oceanography(eds. J. P. RILEY and G. SKIRROW),Vol. 1, pp. 415-496. BREWERP. G. and SPENCERD. W. (197 1) Calorimetric determination of manganese in anoxic waters. Limnol. Oceanogr. 16, 107- 112. BRULANDK. W. (1983) Trace elements in seawater. In ChemicalOceanography(eds. J. P. RILEYand R. -R), Vol. 8, pp. 200-203. Academic Press, London. BURNSR. G. and BURNSV. M. (1979) Manganese oxides. In MarineMinerals(ed. R. G. BURNS),Reviewsin Mineralogy, pp. l-46. Min. Sot. Amer. Publ. 6. EMERSONS., CRANSTONR. E. and LISSP. S. (1979) Redox species in a reducing fjord: equilibrium and kinetic considerations. Deep-Sea Rex 26, 859-878. EMERSON S., KALHORNS., JACOBSL., TEBOB., NEAL~ONK. and ROSSONR. ( 1982) Environmental oxidation of manganese(B): bacterial catalysis. Geochim. Cosmochim. Acta 46, 1073-1079. HEMJ. D. ( 1980) Redox coprecipitation mechanisms of manganese oxides: Particulates in water. In Advances in Chemistry, Series No. 189 (eds. M. C. KAVANAUGHand J. 0. LECKIE),pp. 45-72. American Chemical Society, Washington, D.C. HEM J. D. and LINDC. J. (1983) Nonequilibrium models for predicting forms of precipitated manganese oxides. Geochim. Cosmochim. Acta 47,2037-2046. KALHORNS. and EM!ZRSON S. ( 1984) The oxidation state of manganese in surface sediments of the Pacific Ocean. Geochim. Cosmochim. Acta 48,897-902. KLINKHAMMERG. P. (1980) Early diagenesis in sediments

Appendix:

Dissolved ox‘datio”

spike

time

from the Eastern Equatorial Pacii II. Trace metal results. Earth Planet. Sci. Lett. 49, 8 1- 101. KLINKHAMMERG. P. and BENDERM. L. (1980) The distribution of manganese in the Pacific Ocean. Earth Planet. Sci. Lett. 46, 36 l-384. MARTINJ. H. and KNAUERG. A. (1984) VERTEX: Manganese transport through oxygen minima. Earth Planef. Sci. Lett. 67, 35-47. MORGANJ. J. and STUMMW. (1965) Proceedingsoj’theSetond Conference on Water PollutionResearch, Pergamon, Elmsford, N.Y. MURRAYJ. W., BALISTRIEFU L. S. and PAUL B. (1984) The oxidation state of manganese in marine sediments and ferromanganese nodules. Geochim. Cosmochim. Acta 48, 1237-1247. MURRAYJ. W., DILLARDJ. G., GIOVANOLIR., MOERSH. and STUMMW. (1985) Oxidation of Mn(II): Initial mineralogy, oxidation state and aging. Geochim. Cosmochim. Acta 49,463-470. NEAL+SON K. H. and FORDJ. ( 1980) Surface enhancement of bacterial manganese oxidation: Implications for aquatic environments. Geomicrobiol.J. 2,2 1-37. PIPERD. Z., BASLERJ. R. and BISCHOFF J. L. ( 1984) Oxidation state of marine manganese nodules. Geochim. Cosmochim. Ada 48,2347-2355. ROSSON R. A. and NEALSONK. H. ( 1982) Manganese binding and oxidation by spores of a marine bacillus. J. Bacteriol. 151, 1027-1034. TEBO B. M. and EMERSONS. (1985) The effect of oxygen tension, Mn(I1) concentration and temperature on the microbially catalyzed Mn(I1) oxidation rate in a marie fjord. Appl. Environ. Microbial. 50, 1268- 1293. TEBO B. M,, NEALSONK. H., EMERSONS. and JACOBS S. ( 1984) Microbial mediation of Mn(I1) and Co(B) precipitation at the 02/HrS interfaces in two anoxic fjords. Limnol. Oceanogr. 29, 1247-1258. TIPPINGE., THOMPSOND. W. and DAVIDSONW. (1984) Oxidation products of Mn(I1) in lake waters. Chem. Geol. 44, 359-383. WEISSR. F. (1977) Hydrothermal mangitnese in the deep sea: Scavenging residence time and Mn/He3 relationships. Earth PlanetSci. Lett. 37, 257-262.

manganese concentration and average state of particulate un in expt. A ..-

1 (4/17)

% Mn

(hr)

(PM)

dissolved

0

1.80 1.60 0.28

100 89

CDL CDL

15 2 0

2.80 2.70 :'.a9

1

5.5 8.5 24 50 100 198 2 (q/25)

-1 0 4.2 7.0 27 47

3

(6119)

oxidation

[Mn'"]

0.03 CDL

average

state

2.85

0 0 0

2.86 2.96

0 100

7.96 2.84

91 41 a *

3.30 3.32

3.09

114

0.01

3.35 3.34 3.22

192 216

CDL CDL

3.28

-1 0

24

CDL 1.07 1.81 1.78 0.68

48

0.08

4

3.22

95 36

spike

3.37 3.29 3.25 3.75 3.34 3.44

4 (7123)

time

iM”+5

(hi+)

(IIM)

-1 0

I4

I nn

averageoxidation

dissolved

state

0 100

3.5' 3.37

11.38

96

3.3"

22

10.70

91

3.2h

48 168

9.76 7.34

83 67

3.8 3.e