Chemical and biological reduction of Mn (III) -pyrophosphate

Geochimtca et Cosmochimica Acta, Vol. S9. No. 5. DD. 885-X94. 19% Copynght 0 1995 Ei&vk Science Ltd Printed in the USA. All rights reserved

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Chemical and biological reduction of Mn (III) -pyrophosphate complexes: Potential importance of dissolved Mn (III) as an environmental oxidant

JOEL E. KOSTKA, ’ GEORGE W. LUTHER III,’ and KENNETH H. NEALSON’ ‘Center for Great Lakes Studies, 600 E. Greenfield Avenue, Milwaukee, WI 53204, USA

‘College of Marine Studies, University of Delaware, Lewes. DE 19958, USA

(Received April 22, 1994; accepted in revised form November 15, 1994)

Abstract-Dissolved Mn(III) is a strong oxidant which could play an important role in the biogeochemistry of aquatic environments, but little is known about this form of Mn. Mn( III) was shown to form a stable complex with pyrophosphate which is easily measured by uv-vis spectrophotometry. The Mn ( III) - pyrophosphate complex was produced at concentrations of 5 PM to 10 mM Mn at neutral pH. Inorganic electron donors, Fe( II) and sulfide, abiotically reduced Mn( III) -pyrophosphate in seconds with a stoichiometry of 1: 1 and near 1:2 reductant:Mn (III), respectively. Shewanellu putrefuciens strain MR- 1 catalyzed the reduction of Mn( III) -pyrophosphate with formate or lactate as electron donors. Reduction of Mn( III) catalyzed by MR-1 was inhibited under aerobic conditions but only slightly under anaerobic conditions upon addition of the alternate electron acceptor, nitrate. MR.1 catalyzed reduction was also inhibited by metabolic inhibitors including formaldehyde, tetrachlorosalicylanilide (TCS), carbonyl cyanide m-chlorophenylhydrazone (CCCP), 2-n -heptyl-4-hydroxyquinoline N-oxide (HQNO) , but not antimycin A. When formate or lactate served as electron donor for Mn(II1) reduction, carbon oxidation to CO? was coupled to the respiration of Mn( III). Using the incorporation of ‘H-leucine into the TCA-insoluble fraction of culture extracts, it was shown that Mn( III) reduction was coupled to protein synthesis in MR-1. These data indicate that Mn(II1) complexes may be produced under conditions found in aquatic environments and that the reduction of Mn( III) can be coupled to the cycling of Fe, S, and C.

INTRODUCTION

Little is known about dissolved Mn (III) in aquatic environments. With a reduction potential close to that of molecular oxygen (see equations below), dissolved and solid Mn( III) have the potential to be powerful and important oxidants in marine and freshwater environments. Dissolved Mn(II1) has generally been left out of discussions of the aquatic chemistry of Mn (Stumm and Morgan, 1981), because oxidation states of II and IV were thought to be of greatest importance under conditions found in the natural environment (Morgan, 1967 ) . Mn( III) can undergo reduction or disproportionation according to the following reactions:

Mn”’ + em + Mn”

E” = + 1.5 1 V (Skoog and West, 1982)

MnOOH + 3H’ + c- -+ Mn” + 2H20

E” = + 1.50 V (Stone et al., 1993)

2Mn”’ + 2H20 -+ Mn” + MnOz,s, + 4H+ E” = +0.53 V.

Several studies have alluded to the potential importance of dissolved Mn(II1) as an environmental oxidant in marine and freshwater sediments (Mann and Quastel, 1946; Burdige, 1993; Davison, 1993; Luther et al., 1994) and as an intermediate in the microbial transformation of Mn (Archibald and Fridovich, 1982; Gotffreund et al., 1983, 1985; Ehrlich, 1987; Ghiorse, 1989). Dissolved Mn( III) was also detected as an intermediate in Mn (II) oxidation and aerobic Mn (IV) reduction by bacteria (Gottfreund et al., 1985). Ehrlich ( 1987) suggested that dissolved Mn(II1) may be an intermediate in the anaerobic respiration of Mn( IV) oxides by bacteria, a hy-

pothesis that if proven would have important implications for our understanding of Mn( IV) reduction.

Dissolved complexes of Mn( III) have been studied extensively in soils (Dion and Mann, 1946; Mann and Qaustel, 1946; Heintze and Mann, 1947, 1949; Gottfreund et al., 1985). From extracts of soils with a range of organic matter concentrations and at a variety of pH values, it was suggested that dissolved Mn (III) could arise from direct dissolution of hausmannite or from interaction between pyrolusite or man- ganite with divalent manganese (Heintze and Mann, 1949). Most studies have focused on Mn(III)-pyrophosphate complexes but at least one study was based on complexes made with hydroxycarboxylic acids (Heintze and Mann, 1947). Based on observations that the reverse disproportionation reaction can also take place in solutions of hydroxycarboxylic acid salts and that trivalent manganese is present in soil extracts made with such solutions (Heintze and Mann, 1947), it was suggested that dissolved Mn(II1) may have at least a transient existence in soils ( Heintze and Mann, 1949).

The reduction of solid Mn(II1) by a number of naturally occurring organic compounds has been investigated (Stone and Morgan, 1984a,b; Stone, 1987; Stone et al., 1994; Xyla et al., 1992). Stone and his colleagues have suggested that under anoxic conditions, the reaction of Mn( III) oxides with organic compounds may lead to a substantial portion of the Mn mobilization which has been observed in freshwater and marine sediments. These reactions play an important role in many geochemical transformations including the mobilization and reprecipitation of transition metals in sediments (Stone, 1987). Several studies have shown that solid Mn(III) is produced from the oxidation of Mn( II ) abiotically (Kessick and Morgan, 197.5; Stumm and Giovanoli, 1976; Murray et al.,

885

886 J. E. Kostka. G. W. Luther III. and K. H. Nealson

1985) or catalyzed by Mn-oxidizing bacteria (Adams and

” d

Ghiorse, 1988 )

>

Recently, it was suggested that dissolved complexed

1 d

Mn (III) should be stable or metastable under a range of an-

1) The bacterial reduction of dissolved Mn (III) was coupled

oxic and oxic conditions (Luther et al.. 1994). Further. it was

to carbon oxidation and to protein synthesis.

suggested that Mn(III) complexes might be present in the anoxic waters of the Chesapeake Bay and, therefore, serve as an extremely reactive oxidant in these waters. It is evident that dissolved Mn (III) and its interaction with natural electron donors could play a significant role in the biogeochemistry of

EXPERIMENTAL

freshwater and marine environments. However, if dissolved Mn( III) is to be important as an environmental oxidant, at

Medium

least two conditions must be met: ( 1) it should be easily formed and stable in solution and (2) its reduction should be coupled to the cycling of geochemically important elements such as Fe, S, and C. We implemented a study of the biotic and abiotic transformation of Mn (III ) complexes in the laboratory. In this report, we show that Mn( III)-pyrophosphate complexes are easily produced and stable in solution at concentrations (Mn and ligand) and pH values common to aquatic environments. We used pyrophosphate as it stabilized Mn(III) at the lowest ligand to Mn(II1) ratio and was more easily reduced than Mn( III) complexed to other ligands we tried (such as citrate, malate, pyruvate, tiron, and I .3- bis [ tris (hydroxymethyl)-methylamino] propane). Mn ( III ) - pyrophosphate was chemically reduced by Fe( II) or sulfide, and biologicallv reduced bv Shewanella mrrefaciens (MR-

ilization. The MN3 was adjusted to pH 7.8, filter-sterilized, and di- luted with MN3 where necessary in stoichiometry experiments.

Mn( III)-pyrophosphate complexes were produced from Mn( IV) oxide as well. Mn( IV) oxide was prepared by adding 20 mL 0. I M sodium thiosulfate to 20 ml. 0.1 M potassium permanganate as in Perez-Benito et al. ( 1989). The mixture was centrifuged, washed, and resuspended in distilled water. The Mn(IV) oxide suspension was then added to the Ml medium with 40 mM pyrophosphate and monitored for the production of Mn( III) complexes. A 4 mM soh- tion of Mn( Ill)-pyrophosphate was produced during this partial. reductive dissolution in the Ml medium in 2 days. The Mn(IV) suspension was not characterized and could have contained some Mn(l1).

Bacterial lnoculum

Shewmrllu purr~faciens strain MR-I was used for all biological reduction experiments. This bacterium, originally isolated from the Mn-rich sediments of Lake Oneida, NY (Myers and Nealson, 1988), is a facultatively anaerobic bacterium capable of using a wide variety of electron donors including organic acids, alcohols, and carbohy- drates (Nealson and Myers, 1992). MR.1 was grown aerobically for l-2 days on Luria Broth agar (Maniatis et al., 1982). The cells were rinsed in MN3 2- 3 times, centrifuged in sterile Eppendorf tubes, and resuspended in MN3 or Ml. Inoculum was added to each culture to a final concentration of IO7 cells/ml at time zero.

Mn (III) Reduction Experiments

Following aseptic degassing in sterile erlenmeyer flasks using N2 gas, media were placed immediately into a Coy anaerobic chamber with an atmosphere of 97.5% N,/2.5% Hz for at least 1 hour prior to beginning each experiment. Stock solutions of electron don&s and inoculum cultures of MR.1 were degassed and placed into the anaerobic chamber in the same manner. The electron donor and/or MR- I inoculum was added in the anaerobic chamber and 1 mL was sam- pled from each flask immediately for determination of Mn(III); sub- sequent time points were assessed in the same manner.

A defined, minimal medium (MN3) was used in most of the abiotic and biotic Mn(II1) reduction experiments. MN3 contained per liter of distilled water: 1.19 g ( NH.,)$04. 0.2 g MgS04-7Hz0, 0.057 g CaCI,-2H20, 0.0012 g FeS04-7H20, 0. I mL trace metals supple- ment (Myers and Nealson, I988), 0.1 mL 0. I 15 M NazSe04, Carbon substrates were added to 5- 10 mM. MI medium (Myers and Neal- son, 1988) was used for some experiments, and where appropriate. we have indicated its use and the carbon substrates added. To the Ml medium, yeast extract (Difco) and Bactopeptone (Difco) were added to 0.02 and O.Ol%, respectively. EDTA and HEPES were not used (as in Myers and Nealson, 1988) as both were found to cause abiotic reduction of Mn( III) complexes.

Determination of Mn(II1)

Samples from the reduction experiments were filtered through a disposable nylon filter (Gelman Acrodisc) with a pore size of 0.2 pm and the absorbance was measured on a Beckman DU-64 spectropho- tometer. The absorbance spectrum of the Mn(III) complex was de- termined in the MN3 medium after filtration by scanning from 200 to 800 nm as compared to a blank of distilled, deionized water. Ab- sorbance maxima were observed at 290 and 480 nm. Absorbance at 480 nm was used to monitor the Mn( III)-pyrophosphate concentration unless otherwise indicated.

The Mn (III) -pyrophosphate measurement was calibrated using iodometric titrations nerformed according to Murrav et al. ( 1984). A

Mn(II1) Preparation

The Mn (III ) -pyrophosphate complex was used for the Mn (111) reduction experiments. Sodium pyrophosphate (Aldrich) was added to the MN3 or the Ml medium to a final concentration of 40 mM and the pH was adjusted to 8.2 with concentrated HCI/NaOH. Man- ganese (III) acetate dihydrate (Aldrich) was added to a final concentration of 10 mM and the pH was readjusted to 8.0. Carbon substrate was then added and the medium was filter-sterilized using disposable filterware (Nalgene)

In experiments to determine the stoichiometry of bacterial Mn (III) reduction (Fig. 6, Table 2), the MN3 medium was used in the same manner as described above with the exception that the Mn(III) acetate salt was omitted, and Mn( III)-pyrophosphate was prepared from Mnz03 (Aldrich) Forty grams of MnlOx were added to 500 mL of MN3 containing 40 mM pyrophosphate and the mixture was continuously stirred for 3 days at pH 6.5. After 3 days, the solid was allowed to settle, the mixture was centrifuged, and the supematant was found to contain 9 mM Mn( III)-pyrophosphate after filter-ster-

1 , . ,

standard curve was obtained for the absorbance of Mn(lIl)-pyrophosphate at 480 nm vs. the titrated Mn(II1) concentration ( r2 = 0.9998; Fig. 1) The titrated Mn( III) decreased slightly below that which would be predicted by simple dilution when compared to the total calculated Mn concentration indicating that some Mn( III) was reduced abiotically; however, this represented < 10% of the total Mn added. In conjunction with iodometric titrations of MR-I culture samples, we measured total Mn by the calorimetric method of Brewer and Spencer ( 197 I ).

Protein synthesis was measured using the incorporation of ‘H-leu- tine according to Smith and Azam ( 1992).

Organic acids were measured by ion chromatography using an Aminex anion exchange column (Biorad) and conductivity detection according to Fitzgerald ( 1989).

Carbon dioxide was measured using a Gow Mac gas chromatograph with thermal conductivity detection. Five-mL samples were acidified in IO mL syringes with phosphoric acid and the resulting headspace was injected into the gas chromatograph using He as the mobile phase. Concentration was standardized against a 5 mM bi- carbonate solution.

Dissolved Mn (Ill) as an environmental oxidant 887

Mn3+ + I- 2+ --> 0.5 I, + Mn

0.5 I, + s,o,=--> I- + 0.5 s,o,=

2 r = 0 9998 m = 0 1106

V 2 4 6 a 10 12

n-M Mn(lll)

FIG. 1. Standard curve relating absorbance at 480 nm to titratable Mn(III) obtained by iodometric titration according to Murray et al. (1984). Mn(II1) was complexed with 40 mM pyrophosphate in the MN3 medium.

RESULTS

Formation and Stability of Mn (III) Complexes

Micromolar to millimolar additions of Mn( III) acetate and pyrophosphate (ligand/Mn ratio of 4) resulted in the formation of equimolar amounts of dissolved Mn(III)-pyrophosphate. Once formed, Mn( III)-pyrophosphate was stable under anoxic conditions at pH 5 to 8 (MN3 or Ml ) in the presence of all organic carbon compounds studied (formate, acetate, lactate, pyruvate, and oxalate) for several days. In one experiment at pH 8 in the Ml medium + 10 mM formate, dissolved Mn (III) was stable for over 2 months under anaerobic conditions. The stability of Mn( III)-pyrophosphate was tested and found to be maintained in controls at pH 7-8 in the presence of all amendments to the medium including all poisons and microbial inhibitors used.

Mn( III)-pyrophosphate complexes were measured by spectrophotometry at millimolar concentrations using visible wavelengths and at micromolar concentrations using ultraviolet (uv) wavelengths. Dissolved Mn( III) was complexed and measured from 5 to 100 I_LM at neutral pH in water at a metal/l&and ratio of 4 (data not shown). Maximum absorbance increased with increasing addition of Mn(III), and pyrophosphate complexes showed a maximum absorbance at 2 15 and 260 nm. Mn (III)-pyrophosphate at 2.5 PM appeared to be completely consumed as the peak at 260 nm disappeared upon addition of 100 PM sulfide (data not shown). However, sulfide and other reductants (Fe( IT), hydroxylamine, dithio- nite) produced an absorbance maximum at or below 250 nm (though the absorbance was much lower for sulfide than that observed for Mn(III)-pyrophosphate). In each case, the uv spectra were background corrected using a solution of the ligand alone. We are now pursuing the study of Mn( III) complexed with organic acids in the laboratory.

lodometric titrations were used to show the decrease of Mn(II1) and the absence of Mn( IV) in MR-1 cultures (Fig. 2). Total Mn concentration remained the same while the titrated Mn( III) concentration decreased to zero (Fig. 2). Mn (III) -pyrophosphate complexes were clear with a red tint and no dark precipitate was visible during filtration with or without the addition of MR-1, suggesting that nearly all of the Mn( III) acetate had been complexed in the medium. As expected from the lack of any filterable Mn, no solid Mn( III or IV) was detected by titration. This was supported by the extremely high correlation coefficient observed for the standard curve of titrated Mn( III)-pyrophosphate vs. absorbance at 480 nm (Fig. 1 ) and agreement with total Mn measurements (as well as agreement with total Mn added). A white precipitate was observed during the abiotic and biotic reduction of Mn( III) indicative of a Mn( II) mineral(s) such as rhodochrosite ( MnCO,). We, therefore, can rule out the presence of Mn(IV) in our experiments and we will consider Mn (III) reduction to Mn (II) in the absence of disproportionation.

We also measured the disappearance of lower Mn (III) concentrations in bacterial cultures by absorbance in the ultraviolet range. The peak absorbance shifted with changes in pH and ionic strength. Mn (III) -pyrophosphate formed a relatively broad peak in the MN3 medium from 250 to 320 nm. Measurements at ultraviolet wavelengths were much more sensitive than those in the visible range. The same solution of Mn (III)-pyrophosphate in MN3 gave an absorbance of 0.120 at 480 nm and an absorbance of 2.100 at 290 nm. The ultraviolet peak at 290 nm was observed to decrease from 1.019 to 0.179 in duplicate MR-1 cultures during a 7-hour period. These cultures contained an average of 500 PM Mn( III) when the absorbance was 1.019 at 290 nm. Extinction coefficients for Mn( III)-pyrophosphate in the MN3 medium were 1 .O

a I 52 E - 6

; i

41

2

3 L 0

: ??

4 r;a

-1 04

_.I1-. _~

5 10 15 20 258 30

T!ME (hours)

??Absorbance

0 Mn(lII) - titration with izOA 0 Total Mr. jfcrmoldox~rre:

FIG. 2. Decrease in titratable Mn(II1) as well as absorbance at 480 nm associated with reduction by MR-I in the MN3 medium with 40 mM pyrophosphate and 5 mM formate. Inoculum was IO’ cells/ml.

888 J. E. Kostka, G. W. Luther III, and K. H. Nealson

a

9 L 6

1 Y % 4

2

0

+ 1mM Pyruvate 0

+l. 3mM Fe(U)

v +l, 3mM HS-

/\ = Fe(H), HS-added

1 I 1 I I 0 5 10 15 20 25

TIME (hours)

FIG. 3. Abiotic reduction of Mn(III) by the electron donors: Fe(U), sulfide, and pyruvate in duplicate treatments (MN3 medium at pH 8). Arrows indicate where Fe(I1) or sulfide was added.

X lo* L cm-’ mol-’ and 1.6 X 10’ L cm-’ mol-’ at 480 nm and 290 nm, respectively.

Reduction of Mn ( III)

Chemical reduction of Mn ( III ) -pyrophosphate by Fe( II ) and/or HS occurred rapidly being essentially complete in seconds (Fig. 3; Table 1). One mole of Mn( III) was depleted in the medium for every 1 mole of Fe( II) added and approximately 2 moles Mn (III) were depleted for every 1 mole HS added (Fig. 3; Table 1) No sulfide odor was present and no color formed with ferrozine reagent (a Fe( II)-specific color- imetric reagent) in the Mn(II1) samples after the addition

indicating that the sulfide and Fe( II), respectively, had completely reacted. Acetate, present in the Mn( III) salt, was the only carbon compound present at a significant concentration when testing these inorganic electron donors in the MN3 medium and controls showed repeatedly that no abiotic reduction of Mn (III) occurred in the presence of acetate alone (same as no bacteria controls in Fig. 4). The same results were obtained for chemical reduction by Fe( II) and sullide when these experiments were repeated in the MI medium with IO mM lactate at pH 8.

Kinetic data in Table 1 indicate that the reaction of Mn( III) with Fe(I1) was complete in I.5 s or less. With HS, the reduction of Mn( III) also appeared to be nearly complete in seconds. After addition of Mn(III), no sulfide was detected using the Cline method in the experiments listed in Table I (Cline, 1969; with detection limit of 1 ,uM).

The reduction of Mn( III) by organic compounds occurred at a much slower rate than that observed with inorganic reductants (Fig. 3 ). We observed no significant reduction of 10 mM Mn ( III ) -pyrophosphate in duplicate treatments of 1 mM, 5 mM oxalate, or 1 mM pyruvate in Ml at pH 8 during a 7 day period (all similar to plot of 1 mM pyruvate in Fig. 3 ). A 5 mM pyruvate addition in Ml at pH 8 reduced an average of 1.3 mM Mn( III) in that same period (data not shown).

The reduction of Mn( III) by MR-I was rapid in MN3 and M 1 media (Figs. 4 and 5, respectively). MR-I catalyzed the reduction of Mn( III) at near equal rates with formate or lactate as the electron donor (Fig. 4). We have compared reduction rates in cultures incubated in the absence of Hz (Fig. 4a) to cultures incubated in the presence of H? (Fig. 4b). Incu- bation in the presence of H? enhanced Mn( III) reduction by MR- 1 without carbon substrate added. In duplicate no carbon controls after two days, MR- 1 reduced an average of 0.25 mM without H2 while 1.4 mM Mn(II1) was reduced with H2 (Fig. 4a,b). After 4 days of incubation, an average of 2. I mM Mn(II1) was reduced with H2 and 0.40 mM was reduced with-

Table 1. Stoichiometry resulting from reaction of Mn(lll)-pyrophosphate with inorganic electron donors Fe(ll) and dissoked sulfide. The Ml medium initially contained 10 mM lactate, 9.39 mM Mn(lll), and the pH was 7.95.

Cont. Reaction mM Mn(lll) A Mn(lll) Added Time (set) remaining

Fe(ll)- 2.45 mM

4.79 mM

9.15 mM

15 60

60

15 60

6.97 7.04

4.72

0.41 0.59

2.42 2.35

4.67

6.98 6.60

Fe/Mn Ratio 1.0 1.0

1.0

1.0 1.0

HS-- 1.52 mM 60

300

2.97 mM 15 60 300

4.36 mM 15 60

6.89 2.56 6.76 2.63

4.85 4.54 4.47 4.92 4.17 5.22

2.84 6.55 2.36 7.03

S/Mn Ratio 0.6 0.6

0.7 0.6 0.6

0.7 0.6

Dissolved Mn(lII) as an environmental oxidant 889

c

y 6 t

Y 4 I

2

0

09 10

‘1, \ k

%; MR-1 + 5 mM lactate, 4 MR-1 + 5 mM formate

‘?k---_ I 1 / 1

0 10 20 30 40 50 60

t

c

TIME (hours)

7 / , I I I No bacteria

n -r, ” s 0

MR-1 + no carbon

MR-1 + 5 mM lactate, MR-1 + 5 mM formate

J 0 10 20 30 40 50 60

TIME (hours)

FIG. 4. Reduction of Mn(lI1) by MR-I in duplicate cultures in the MN3 medium. (a) sealed cultures without Hz added (deaerated and charged with argon), and (b) incubated in the anaerobic hood in an Hz-containing atmosphere. Inoculum was IO’ cells/ml.

out Hz in these same cultures (data not shown). Rates with

formate or lactate added were similar in the presence and absence of H2 (Fig. 4a,b). Reduction of Mn( III) by MR-1 was optimal at pH 7 to 8 with rates much lower at pH 6, and almost no reduction was observed in two days at pH 5 (data not shown )

Experiments with poisons and metabolic inhibitors suggest that reduction of Mn( III) is linked to respiration in MR-1. The protonophores m-chlorophenylhydrazone (CCCP) and tetrachlorosalicylanilide (TCS), inhibited the reduction of Mn(II1) by 50 and lOO%, respectively (Fig. 5a), while electron transport inhibitors gave contrasting results. Antimycin A did not inhibit Mn( III) reduction under the conditions tested while 2-n-heptyl-4-hydroxyquinoline n-oxide (HQNO) inhibited Mn (III) reduction by 92% (Fig. 5b). Formaldehyde also completely abolished Mn (III) reduction by MR-1.

Under aerobic conditions, Mn( III) reduction by MR- I was 90% inhibited (Fig. Sa), while with 5 mM nitrate (and 9 mM Mn(II1)) added, only a 16% inhibition of Mn( III) reduction was observed (Fig. 5a). Further, when equal 2 mM amounts of nitrate and Mn( III) were added under similar culture conditions, virtually no inhibition of Mn( III) reduction was observed (data not shown).

Mn( III) reduction was coupled to carbon oxidation in MR- 1 (Fig. 6). A linear increase in CO2 was produced in MR-1 cultures with increasing additions of Mn( III) and formate (r* = 0.977) or lactate ( r2 = 0.996) as the carbon substrate (Fig. 6). Measuring the decrease in organic acid concentration, we found Mn/C ratios of 1.8: 1 and 4.7: 1 for formate and lactate

(4 lo

a

6 9 E

Y 4

z

2

0

r I I I I I_/?__ + Formaldehyde,

+ TCS

\r\

‘1 -w Unmhlbked’

4

4

_J

0 5 10 15 20 25 30 35

TIME (hours)

Umnhiblted

+Ant 2 -

0-L-I 0 5 10 15 20 25

TIME (hours)

FIG. 5. Reduction of Mn(lII) by MR.1 in the MI medium + 10 mM lactate (with yeast extract and peptone added) in duplicate treatments including (a) +0.2% formaldehyde, 100 PM CCCP, 100 PM TCS, 5 mM nitrate, and aerobic treatments, (b) 50 PM HQNO, 100 PM antimycin A, and uninhibited treatments (in both a and b). In- oculum was 10’ cells/ml. Controls showed no abiotic reduction of Mn(II1) by equal concentrations of microbial inhibitors in the absence of cells.

890 J. E. Kostka. G. W. Luther 111. and K. H. Nealson

I

I / / _ i 0 2 4 6 8

Mn(lll) Reduced (mM)

FIG. 6. The CO* produced from the oxidation of formate and the oxidation of lactate with increasing additions of Mn(II1) in duplicate MR.1 cultures with + 5 mM carbon substrate added. No carbon and no electron acceptor controls were carried out in duplicate to measure the background of CO? produced and Mn(II1) reduced. All cultures were grown in the absence of Hz in MN3.

oxidation, respectively (Table 2). Using the CO, produced (data shown in Fig. 6), mean stoichiometries were 1.8 and 3.6 for formate and lactate oxidation, respectively. Mn( III) concentrations were adjusted for the Mn(II1) reduced in no carbon and no bacteria controls. Carbon concentrations were adjusted for the amount of carbon oxidized in no electron acceptor controls.

Protein synthesis was coupled to Mn( III) reduction using the incorporation of 3H-leucine. Figure 7 shows that duplicate MR-1 cultures with Mn(II1) added incorporated ‘H-leucine over controls to which no electron acceptor was added. A

50000

40000

30000

20000

10000

0

/v\ + 10 mM Mn(II1)

No electron acceptor 0

1

_J

0 4 a 12 16 20 24

TIME (hours)

FIG. 7. The incorporation of ‘H-leucine into the TCA-insoluble fraction of duplicate culture samples with Mn(II1) (triangles) and with no electron acceptor added (circles). MR.1 was cultured in the Ml medium with 10 mM lactate added as the carbon substrate.

decrease in Mn (III) corresponded to the incorporation of ‘H- leucine in duplicate cultures of MR-1.

DISCUSSION

Our interest in dissolved Mn (III) was to explore its potential as an environmental oxidant by studying its stability under anoxic conditions and by studying its reactivity during chemical and microbiological reduction. Dissolved Mn (III) was shown to react with organic (pyruvate) and inorganic electron donors (Fe( II), H2S) and was coupled to the oxidation of formate or lactate by the Mn-reducing bacterium, Shewanella putrefuciens (MR-1). These experiments show that stable Mn (III) complexes can form and be measured at micromolar to millimolar concentrations, and these data support previous

Table 2. Stoichiometry of Mn(lll)-pyrophosphate reduction catalyzed by MR-1 including organic acid measurements, Mn(lll) measurements, and Mn/C ratios. a) with formate as the electron donor, b) with lactate as the electron donor. Each value is an average from duplicate cultures. These cultures were grown in an Hz-containing atmosphere. Mn(lll) and organic acid concentrations were corrected using no carbon and no electron acceptor controls.

4

0 hrs

16 hrs

b)

0 hrs

16 hi-s

Formate (mM)

6.64

3.64

Lactate (mM)

5.74

4.72

Mn(lll) Mn/C (mM)

7.89

2.66 1.6/l

Mn(lll) Mn/C (mM)

7.90

3.30 4.7/l

Dissolved Mn(II1) as an environmental oxidant 891

work which suggests that Mn( III) complexes may well form under conditions commonly found in marine and freshwater environments (Heintze and Mann, 1947; Popp et al., 1990; Luther et al., 1994). Dissolved Mn (III) could be involved in the rapid geochemical cycling of Fe and S, or as an electron acceptor during the dissimilatory reduction of Mn with the concomitant oxidation of organic matter.

Chemical Reduction of Mn (III)

Our data show a 1: 1 ratio between Fe( II) added and Mn( III) reduced and a ratio approaching 1:2 was observed for HS _ added and Mn (III) reduced (Fig. 3; Table 1) We suggest the abiotic reduction of Mn( III) proceeds with Fe( II) and HS as inorganic electron donors according to the following reactions:

Mn”’ + Fe” -+ Mn” + Fe”’

2Mn”’ + HS- -+ 2Mn” + S” + Ht.

Mn (III), Fe( II), and Mn (II) were all likely partially or completely complexed to pyrophosphate in our experiments. Mn( II) also formed a white solid which dissolved in acid that is probably a carbonate solid such as rhodochrosite (MnCO,). We could not detect any Fe( II) in the medium after adding it to the Mn( III), further suggesting that the reaction was complete. While we did not confirm the sulfide oxidation reaction through the measurement of S”, the stoichiometry of the re- actants suggests the reaction shown above. The observed ratio of sulfide to Mn(II1) (Table I), though approaching 0.5, was slightly higher at 0.6 to 0.7. Since we did not check our sulfide for purity, the ratio could be higher than that predicted in the balanced reaction above due to impurities or some of the sulfide may have been lost during transfer to the Mn (III).

Detailed rate studies were not done, but we observed the loss of red color (Mn( III)) almost immediately upon addition of Fe( II) or sulfide, and the Mn( III) concentration was measured within seconds of the addition of electron donor (Fig. 3; Table 1). The reduction of dissolved Mn ( III) (Table 1) is 100 to 1000 times as rapid as those rates reported for the reduction of solid Mn(IV) by Fe(U) (Burdige et al., 1992) and approximately equal to the reduction rate of solid Mn( IV) reported for dissolved sulfide (Burdige and Nealson, 1986 )

Studies by Burdige and co-workers (Burdige et al., 1992; Bur- dige and Nealson, 1986) used much higher electron acceptor (solid Mn( IV)) concentrations and one-half to one-tenth the concentration of electron donor (sulfide, Fe(I1)) as in this study. Such rapid reactions could be responsible for the re- cycling of oxidized Fe and S in sediments and in waters with oxic/anoxic interfaces like the Black Sea (Luther et al., 1991), although more detailed kinetic comparisons are needed. With dissolved Mn(III), there would be no surface area constraints to limit the reaction rate as there would be with solid Mn (IV ) as the oxidant.

Bacterial Reduction of Mn (III)

Ehrlich (1987) mentioned the possible importance of Mn( III) as an intermediate produced during the anaerobic respiration of Mn (IV) by bacteria, and he questioned whether

the reduction of Mn( IV) was a one-step, two-electron process or a two-step process according to the following reactions:

One-step- MnO, + 4H’ + 2e- --t Mn” + 2H20

Two-step-

MnO, + 4H + + e --t Mn”’ complex + 2H20

Mn”’ complex + c + Mn”.

The data of Gottfreund et al. ( 1983, 1985 ) suggest that dissolved Mn(II1) may be produced as an intermediate during aerobic Mn (IV) reduction, while our data show that bacteria can couple the anaerobic respiration of Mn(II1) to the oxidation of organic matter. Inhibition of Mn(II1) reduction by protonophores CCCP and TCS (Fig. 5a) infers that MR-1 couples Mn(II1) reduction to energy generation through the production of a proton-motive force. This coupled with the almost complete inhibition of Mn (III ) reduction under aerobic conditions (Fig. 5a) points to true anaerobic respiration in MR- 1. Apparently, the reduction of Mn (III ) involving only a single electron transfer is capable of producing energy for the bacterial cell, and the two-step process suggested by Ehr- lich (1987) is at least possible. This is not surprising as the reduction of Fe( III) to Fe( II) by Shewanrflu putrefaciens has been studied extensively (Lovley, 199 1; Nealson and Myers, 1992).

The lack of inhibition of Mn( III) reduction by antimycin A (Fig. 5b) is in contrast to results for Mn( IV) reduction by MR-1 which was inhibited by 67% (Myers and Nealson, 1988). Also, proton translocation linked to Mn( IV) reduction was inhibited by antimycin A in a single experiment (Myers and Nealson, 1990). However, Mn( IV) reduction was not inhibited by antimycin A in Bacillus 29 (Ghiorse, 1988). In- hibition by antimycin A is associated with the involvement of cytochrome b. Lack of inhibition suggests Mn (III) reduction follows a different electron transport pathway than Mn(IV) reduction in MR- I.

Mn( III) reduction, which is almost completely stopped under aerobic conditions, proceeds relatively unabated in the presence of nitrate (Fig. 5a) It appears that the energy de- rived from Mn( III) reduction by bacteria lies somewhere between that of O? and nitrate reduction which agrees with the thermodynamics of the predicted half reactions (Froelich et al., 1979) and further suggests that Mn reduction may occur between zones of oxygen and nitrate reduction in sediments and anoxic basins.

MR-I couples the oxidation of organic carbon to Mn( III) reduction (Table 2; Fig. 6). Oxidation of formate and lactate follow predictable stoichiometries based on the depletion of organic acids and the production of CO?. Our results show that 1.8 moles of Mn( III) are reduced for every 1 mole of formate oxidized by MR- 1 and 3.6-4.7 moles of Mn( III) are reduced for every 1 mole of lactate oxidized according to the following equations:

2Mn”’ + HCOO + HZ0 -+ 2Mn” + HCO? + 2H +

4Mn”’ + CH,CHOHCOO- + 2H20 --t 4Mn”

+ CHKOO + HCO, + 5H’

2Mn”’ + HZ + 2Mn” + 2H’.

892 J. E. Kostka. G. W. Luther 111, and K. H. Nealson

We measured acetate production in one experiment and found an increase of 300 /IM which corresponded to a decrease of 1 mM lactate. As predicted by the balanced equation above, equimolar amounts of acetate should be produced from lactate oxidation coupled to Mn(II1) reduction. MR-1 may have incorporated acetate as a carbon source as has been shown previously (Myers and Nealson, 1990) which would explain the lower acetate concentrations we measured. Further detailed studies are needed on the carbon metabolism of MR- 1 coupled to Mn( III) respiration.

Our data suggest that MR-1 is capable of coupling the oxidation of H2 to the reduction of Mn( III) (Fig. 4; see Results section). Higher Mn( III) reduction rates were observed in the presence of H2 than in cultures grown in Nz or Ar alone. Also, the Mn/C ratio observed for lactate oxidation in the presence of H2 (Table 2) may be higher than the Mn/C ratio of cultures which were not exposed to HZ (Fig. 6) due to this coupling.

Whether or not MR-1 couples reduction of Mn(II1) to growth will require further experimentation. Our prelimi- nary data show that MR-1 does synthesize protein when respiring Mn( III) (Fig. 7); however, we have not been able to measure an increase in cell number. We suggest that Mn(III), a highly reactive oxidant, may be toxic to cell growth but not to respiration at the concentrations used in this study. This toxicity does not seem to be due to Mn(I1) concentrations as MR-1 is not inhibited in the same way during Mn( IV) reduction in the presence of Mn (II) (My- ers and Nealson, 1988).

Biogeochemistry of Mn (III)

Dissolved Mn( III) has been largely ignored as a significant environmental oxidant, probably because stability constants for inorganic species encountered in natural waters (phosphate, sulfate) are too low for those species to be significant ligands of oxidized Mn, and the occurrence of organic complexing agents capable of solubilizing Mn(II1) was simply not known (A. Stone, pers. commun.). However, given the biological and chemical results described here, Mn( III), both dissolved and solid forms, could have significant effects on the biogeochemistry of freshwater and marine environments. Mn( III) could be responsible for the rapid geochemical cycling of Fe, S, and C as well as serve as a substrate for Mn- reducing bacteria. Soluble Mn (III) complexes will undoubt- edly be reduced by inorganic species (e.g., Fe(II), HS ), some organic compounds, and bacteria. Malonate has been shown to reduce dissolved Mn (III) (C. Burkhard and G. W. Luther III, pers. commun.) Regardless of whether the reduction of Mn(II1) complexes is fast or slow, their existence is very important because they represent a means of transporting a strong oxidant via aqueous diffusion and dispersion.

More questions than answers remain regarding the potential significance of dissolved Mn( III). Since monodentate inorganic complexes are not feasible, Mn( III) is likely to be complexed by organic complexing agents with sufficiently large complex formation constants. Therefore, bacteria may well be involved as they are capable of producing chelates of Mn( III). The data so far indicate that Mn( III) complexes are likely to be produced from the dissolution of solid Mn( III),

solid Mn( IV), solid Mn( II&IV), or from the oxidation of complexed Mn(I1) in aquatic environments. Contrary to the suggestions of Heintz and Mann ( 1949), we suggest that the back reaction of disproportionation would produce solid over dissolved Mn (III) due to the instability of most Mn( III) complexes. In the simple model below, we show the oxidation states likely to occur in nature, all interconvertible, and each with a solid (s) and complexed (c) form.

1 Mnf-Mn; 1

The central process to be explored in the future is the complexation of Mn (III) (and all oxidation states for that matter) by organic complexing agents and interaction of these complexes with solid phase Mn. Mn(II1) and Mn(IV) complexes have been shown as a soluble cluster in the photosystem center of plants (Harriman, 1979) and have been suggested to occur in aquatic environments (Chiswell and Mokhtar, 1986). Also, a soluble form of colloidal Mn (IV) has been produced in the lab (Perez-Be- nito et al., 1989). Reductive and nonreductive transformations of all of these Mn species in the environment are likely to be mediated by bacteria and plants.

Field studies of Mn geochemistry, such as those of Stauffer ( 1986), Aguilar and Nealson ( 1994), and others, have tra- ditionally been done by filtration, assuming that dissolved Mn is equivalent to reduced Mn or Mn (II), and conversely, that particulate Mn is composed mainly of Mn (IV). If a dissolved oxidized form of Mn (Mn(II1)) exists, it would be of great importance to the biogeochemistry of Mn and it should be sought out in the field. Also, the importance of the cycling of solid Mn( III) in sediment geochemistry should be considered further. Solid forms of Mn( III) are fairly well characterized in the laboratory. Hausmannite (Mn,O,) and feitknechtite (p- MnOOH) have been identified as primary products of Mn (II) oxidation (Kessick and Morgan, 1975; Hem and Lind, 1983; Murray et al., 1985) and these Mn( III) oxides have been found to be reduced by a number of organic compounds found in nature (Stone and Morgan, 1984a,b; Stone, 1987; Stone et al., 1994; Xyla et al., 1992). The fact that bacteria are capable of reducing Mn(II1) and coupling it to the oxidation of organic matter adds yet another complexity. So far, Mn(II1) oxides have been left out of studies which focus on the inter- actions between Mn oxide mineralogy and microbial Mn reduction (Burdige et al., 1992).

Evidence for complexed Mn (III) has been presented for agricultural soils (Mann and Quastel, 1946) where Mn reacts rapidly with organic root exudates (Jauregui and Reisenauer, 1982), and it has been hypothesized that Mn(II1) may play a pivotal role in nutrient cycling. Manganese has been found to dominate the sedimentary geochemistry of some lake environments (Dean et al., 1981) and marine environments (Canfield et al., 1993). These environments will provide a good starting point to search for dissolved complexed Mn( III) and to explore its potential importance as an environmental oxidant.

Dissolved Mn(lll) as an environmental oxidant 893

1)

2)

3)

4)

CONCLUSIONS

At Mn concentrations and pH values common to aquatic environments (and at a ligand/Mn ratio of 4/ 1 ), Mn( III) forms complexes with pyrophosphate. These Mn (III) complexes can be kept stable in solution and can be measured easily using uv/vis spectrophotometry and iodometric titrations. Chemically, Mn( III) -pyrophosphate is reduced almost in- stantaneously by Fe( II) and sulfide, while the reduction is slower or nonexistent with pyruvate and oxalate as the electron donor under the conditions used in our experiments. Bacteria isolated from Mn-rich lake sediments respire Mn (III) and that respiration is coupled to carbon oxidation and protein synthesis.

Acknowledgmenfs-We would like to thank Russell Cuhel for help with the leucine uptake experiments and Jim Waples for his aid with the CO? measurements. We are grateful for the careful reviews of David Burdige, Jim Morgan, Alan Stone, and one anonymous re- viewer. This work was supported by a NSF Postdoctoral Fellowship in Marine Biotechnology (OCE-9212460 to JEK), a NSF grant to GWL (OCE-93 13020), and a NASA grant to KHN (#NAGW3270).

Editorial handling: G. R. Holdren Jr.

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Chemical and biological reduction of Mn (III) -pyrophosphate