Related Electrochemical Characteristics of Microbial Metabolism and

Mar 6, 2000 - half-cell reactions,2 which describe the corrosion of iron in acidic ... Pourbaix diagram for iron at 25 °C. The activities of all ...
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Ind. Eng. Chem. Res. 2000, 39, 575-582

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APPLIED CHEMISTRY Related Electrochemical Characteristics of Microbial Metabolism and Iron Corrosion Denny A. Jones* Metallurgical and Materials Engineering, University of Nevada, Reno, Reno, Nevada 89557

Penny S. Amy Biological Sciences, University of Nevada, Las Vegas, Las Vegas, Nevada 89154

Both bacterial growth and corrosion involve electrochemical (charge-transfer) half-cell reactions which have been conveniently compared on the potential-pH (Pourbaix) diagram in this paper. We have tabulated bacteria which mediate or facilitate various pairs of these half-cell reactions in otherwise kinetically unfavorable electrochemical cells. The tabulated electrochemical cells cover the range of microbial physiology from oxidizing (aerobic) to reducing (anaerobic) conditions and from acidic to neutral and alkaline pH on the Pourbaix diagram. Correlations emerging from Pourbaix plots of microbially affected half-cell reactions are discussed in relation to general microbial ecology and specifically to their influence on the corrosion of iron and carbon steel. Introduction Multitudes of bacteria are ubiquitous in common domestic, cooling, cleaning, and chemical process waters at ambient and near-ambient temperatures; they await only the necessary conditions to metabolize and grow. Each bacterial strain is genetically adapted to facilitate or mediate one or a few related electrochemical cells. The corresponding cell reactions are normally quite slow without microbial mediation and occur only in specific conditions of pH and electrochemical potential. This potential measures the oxidizing “power”, or the extent of aeration/deaeration, in the solution. Figure 11 illustrates areas on the potential-pH (Pourbaix) diagram where some selected major types of bacteria are active. These bacteria are important in microbial ecology, geochemistry, water treatment, and corrosion. That version of the diagram in Figure 1 was chosen to show hydroxides, Fe(OH)2, and Fe(OH)3 as the stable solid phases most appropriate for aqueous solutions. The similar version showing the corresponding oxides, Fe2O3 and Fe3O4, as the stable phases does not appreciably affect the interpretations in this paper. Aqueous corrosion also progresses by similar electrochemical cells without bacteria and Pourbaix diagrams are widely employed to describe electrochemical equilibria in aqueous corrosion. In this paper, we review and compare the electrochemical characteristics of corrosion and bacterial metabolism and briefly relate them to the mechanism of microbiologically influenced corrosion (MIC) of iron and low-carbon steel. Electrochemical Corrosion of Iron The following list shows the usual electrochemical half-cell reactions,2 which describe the corrosion of iron * To whom correspondence should be addressed. Tel.: (775) 784-6021. Fax: (775) 327-5059. E-mail: [email protected].

in acidic, neutral, and alkaline pH:

2H+ + 2e- ) H2 eH+/H2 ) 0.000 - 0.059 pH

(1)

O2 + 2H2O + 4e- ) 4OHeO2/H2O ) 1.229 - 0.059 pH

(2)

Fe ) Fe2+ + 2eeFe/Fe2+ ) -0.440 + 0.0295 log(Fe2+)

(3)

Fe + 2H2O ) Fe(OH)2 + 2H+ + 2eeFe/Fe(OH)2 ) -0.0470 - 0.0591 pH

(4)

Fe2+ + 2H2O ) Fe(OH)2 + 2H+ pH ) 6.65-0.5 log(Fe2+)

(5) +

Fe(OH)2 + H2O ) Fe(OH)3 + H + e

-

eFe(OH)2/Fe(OH)3 ) 0.271 - 0.059 pH

(6)

Fe2+ + 3H2O ) Fe(OH)3 + 3H+ + eeFe2+/Fe(OH)3 ) 1.057 - 0.1773 pH - 0.0591 log(Fe2+) (7) Fe3+ + 3H2O ) Fe(OH)3 + 3H+ pH ) 1.613 - (1/3)log(Fe3+)

(8)

Fe2+ ) Fe3+ + eeFe2+/Fe3+ ) 0.771

10.1021/ie990516v CCC: $19.00 © 2000 American Chemical Society Published on Web 03/06/2000

(9)

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Figure 1. Pourbaix diagram showing areas of activity for various classes of bacteria. IB: iron bacteria. TB: thiobacteria. DNB: denitrifying bacteria. SRB: sulfate-reducing bacteria. HAB: heterotrophic anaerobic bacteria.

Figure 3. Pourbaix diagram for iron with additional superimposed lines representing reactions significant in MIC. Numbers in ( ) correspond to similarly designated reactions in the text. Activities of all dissolved species assumed to be 10-2 mol/L.

MIC. Increasing or more noble positive potential on the diagram measures the activity of dissolved oxidizers (often dissolved oxygen). Stated another way, decreasing dissolved oxygen (deaeration) is reflected by potentials decreasing below the line labeled for (2) at any given pH, assuming no other dissolved oxidizers are present. Further features of the diagram, corresponding to more alkaline reactions, are omitted above pH 11 where microbial growth is limited. Figure 2 describes electrochemical equilibria of iron or low-carbon steel corrosion in aqueous systems. An electrochemical corrosion cell is formed when a noble (positive) half-cell reaction is cathodically reduced (electron acceptor) in combination with another more active half-cell reaction which is simultaneously oxidized (electron donor). For example, summing (1), reduction, and (3), oxidation, produces the cell reaction,

Fe + 2H+ ) Fe2+ + H2 Figure 2. Pourbaix diagram for iron. Numbers in ( ) correspond to similarly designated reactions in the text. Activities for all dissolved species assumed to be 10-2 mol/L.

The equilibrium for each reaction is described by its half-cell electrode potential, e, according to the corresponding Nernst equation,2 which follows each reaction. The Nernst equations for (1)-(9) result in linear plots of potential versus pH in Figure 2, which defines the Pourbaix diagram for iron at 25 °C. The activities of all dissolved ions (enclosed by parentheses in the Nernst equations) are taken somewhat arbitrarily as 10-2 mol/L in Figure 2 for comparison with assumed conditions of

for corrosion of iron in deaerated acid solutions. Only the cathodic reactions for the reduction of H+ (1) and dissolved oxygen (2) are normally included in the Pourbaix diagram because they are the only ones with adequate concentration (more correctly activity) and room-temperature rates to affect corrosion. However, reaction (9), the reduction of Fe3+ to Fe2+, could be a factor at 10-2 mol/L in acidic solutions and is included in Figure 2. All the rest, (3)-(8), become anodic oxidation reactions for iron corrosion when combined with (1), (2), and/or (9) in electrochemical corrosion cells.

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6CO2 + 24H+ + 24e- ) C6H12O6 + 6H2O

Bacterial Mediation of Electrochemical Reactions Figure 3 superimposes Nernst-equation plots for a number of additional half-cell electrode reactions, (10)(20), on the Fe-Pourbaix diagram of Figure 2. They are listed below with their corresponding Nernst equations, in approximate order of descending potential, and were selected for their known importance in general microbial ecology3,4 and/or MIC.5 The Pourbaix Atlas2 is the source for the Nernst equations in (10)-(18), while those in (19) and (20) were derived from thermodynamic data listed by Lindsay.6 Activities of all dissolved species are again assumed to be 10-2 mol/L, as in Figure 2.

2NO3- + 12H+ + 10e- ) N2 + 6H2O e ) 1.25 - 0.071 pH + 0.006 log(NO3-)/PN2

(10)

MnO2 + 4H+ + 2e- ) Mn2+ + 2H2O e ) 1.23 - 0.118 pH - 0.03 log(Mn2+)

(11)

NO3- + 2H+ + 2e- ) NO2- + H2O e ) 0.835 - 0.059 pH

(12)

NO2- + 8H+ + 6e- ) NH4+ + 2H2O e ) 0.897 - 0.079 pH

(13)

HSO4- + 7H+ + 6e- ) S0 + 4H2O e ) 0.339 - 0.069 pH + 0.01 log(HSO4-) (14) SO42- + 8H+ + 6e- ) S0 + 4H2O e ) 0.357 - 0.079 pH + 0.01 log(SO42-)

(15)

SO42- + 10H+ + 8e- ) H2S + 4H2O e ) 0.311 - 0.074 pH + 0.007 log(SO42-)/PH2S (16) S0 + 2H+ + 2e- ) H2S e ) 0.142 - 0.059 pH - 0.03 log(H2S)

(17)

N2 + 8H+ + 6e- ) 2NH4+ e ) 0.275 - 0.079 pH + 0.01 log PN2/(NH4+)2 (18) CO2 + 8H+ + 8e- ) CH4 + 2H2O e ) 0.169 - 0.059 pH

(19)

e ) 0.015 - 0.059 pH - 0.002 log PCO26/(C6H12O6) (20) Specific bacteria have genetically evolved the metabolic capability to mediate (enhance the rate of) electrochemical cell reactions formed from pairs of the listed half-cell reactions (1)-(20). Abiotic reaction rates are often rather low at room temperature in dilute concentrations, despite the rather high potential differences between the half-cell reactions. For example, there is little or no abiotic ambient-temperature reaction in the electrochemical cell between reduction of dissolved oxygen (2) and oxidation of glucose (20), but a multitude of aerobic bacteria beneficially use the large electrochemical half-cell potential difference between the two for energy in the growth of bacterial cells and incorporate some forms of oxidized carbon into the structure of new cells. The large-energy, single-step reaction (20) does not represent the mechanism, only the overall equilibrium between glucose and its oxidation product, CO2. Reaction 20 is taken here to be representative of most organic materials which can be oxidized to CO2 with help from bacteria. The oxidation mechanism is well-known to involve a complex series of biochemical steps, including the consecutive pathways of glycolysis, the tricarboxylic acid cycle, and finally oxidative phosphorylation.7 The mechanism is sufficiently versatile to allow a diversity of bacteria to efficiently metabolize a huge variety of organic compounds and mixtures. Table 1 is a partial list of those bacteria which mediate pairs of the various half-cell reactions ((1)(20)). Oxidation reactions are displayed horizontally across the top of Table 1, showing the reduced form being oxidized to a higher state. Conversely, the corresponding reduction reactions are displayed vertically, along with the half-cell electrode potential at pH 4 chosen as representative for biological mediation. Bacteria which mediate a given cell reaction may be found at the intersection of the column corresponding to the anodic oxidation half-cell reaction with the row for the cathodic reduction half-cell reaction. Any of the halfcell reactions ((1)-(20)) can be anodically oxidized or cathodically reduced, in principle, depending on whether its half-cell potential is active or noble, respectively, to that of its partner in the electrochemical cell. Unless otherwise noted, the bacterial species listed in Table 1 were largely derived from Atlas and Bartha,3 Ehrlich,4 Geesey,5 and Laskin et al.,8 who include most of the original references. Table 1 omits (15), reduction of HSO4- to S0, which occurs largely at pH < 1 where biomediated reactions are rare. Figure 3 and Table 1 provide a fairly comprehensive, but certainly not exhaustive, sample of the many electrochemical cell reactions subject to microbial mediation. Any blank spaces in Table 1 do not necessarily mean that there are no bacteria which mediate the electrochemical cell between the corresponding pairs of half-cell reactions, only that they are sufficiently rare that they have not been mentioned in the literature surveys3,4,5,8 consulted. The significance of the various biomediated cells in Table 1 is reviewed in the following section with emphasis on their effects on corrosion of iron and carbon steel. The review is facilitated by Figure 4, which compares specific areas of microbial activity shown in

a The “*” indicates the following: numbers in parentheses ( ) are the corresponding half-cell reactions in the text. Numbers in parentheses with prefix (A) refer to page numbers in Atlas and Bartha,2 (E) to page numbers in Ehrlich,3 (G) to page numbers in Geesey,4 and (L) to page numbers in Laskin et al.7 Numbers in brackets [ ] refer to the appended list of references.

Table 1. Microorganisms Mediating Electrochemical Reactionsa

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Figure 4. Comparison of the reaction lines from Figure 3 with corresponding areas on the Pourbaix diagram of Figure 1 for (a) ironreducing bacteria (IRB) and thiobacteria (TB), (b) sulfate-reducing bacteria (SRB), (c) nitrogen bacteria including denitrifying bacteria (DNB), and (d) heterotrophic anaerobic bacteria (HAB).

Figure 1 with the relevant electrochemical reactions selected from Figure 3. Discussion Many aerobic, heterotrophic bacteria, for example, Pseudomonas, facilitate the oxidation of organic material such as glucose, C6H12O6, (20), to produce H2O and CO2, while reducing dissolved oxygen, the “terminal electron acceptor” or oxidizer (2). The large cell potential evident in Figure 3 between the lines for (2) and (20) provides the driving force for cell growth. The bacteria incorporate part of the organic carbon into the structure of existing and new cells during growth and reproduction, respectively. Many aerobic heterotrophic bacteria are known for copious production of exopolymeric substances (EPS) or slime coatings to protect themselves

and others from changes in the environment. Although most bacteria have some of their own slime-forming capabilities, they may prefer to coexist cooperatively (symbiotically) with other more efficient slime producers to create a protective “biofilm”. The biofilm and associated EPS may be incorporated with iron oxides to produce bulk surface deposits or localized tubercles under which occluded electrochemical corrosion cells can prosper. Some oxides, especially magnetite, Fe3O4, may have sufficient electronic conductivity to allow their surfaces to serve as cathodes. Anaerobic heterotrophic bacteria also oxidize organic matter by (20) but utilize different oxidizers (electron acceptors), which function at somewhat lower half-cell electrode potentials in place of or along with dissolved oxygen. Some species, including Thiosphaera pantotro-

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pha, can reduce NO3- heterotrophically (10) in the absence of or together with dissolved oxygen, probably because the potentials of the oxygen (2) and nitrate (10) reductions are not far apart (Figure 3). Other bacteria do not utilize inorganic electron acceptors such as oxygen and nitrate, but reduce organic molecules producing organic acids, carbon dioxide, and hydrogen gas. Fermentative bacteria, such as Escherichia coli, oxidize organic matter for both carbon and energy (20) while reducing organic molecules to form organic acids and other reduced fermentation products. For example, lactobacilli ferment carbohydrates and produce organic acids such as lactic and acetic acids at very low pH values (pH 3-4) and are used in the food industry to produce yogurt and sauerkraut. Some products of the anaerobic fermentation process play a direct role in MIC, for example, hydrogen sulfide, hydrogen, and acids. The importance of electrochemical potential (charge transfer) over chemical species (chemistry) is apparent. Heterotrophic bacteria are available to mediate the reduction of most half-cell reduction reactions listed in Table 1, testifying to the convenience and necessity of organic matter as a nutrient for aerobic bacteria. Otherwise, all carbon would remain bound in the organic state and life would cease. Autotrophic bacteria, which mediate the oxidation of inorganic matter, must derive carbon for cell growth by extraction (fixation) of CO2 from ambient air or water. Aerobic, iron-oxidizing bacteria, for example, Thiobacillus ferrooxidans, T. prosperus, and Leptospirillum ferrooxidans, thrive in acidic environments (acidophiles) where ferrous ions are more stable in oxidizing conditions at high electrochemical potentials (Figure 4a). They use the energy from the oxidation of the ferrous ion, Fe2+, to the ferric state in Fe(OH)3 (7) with simultaneous reduction of dissolved oxygen (2). The oxidized Fe(OH)3 is much less soluble than the parent Fe2+, and insoluble deposits (tubercles) of Fe(OH)3, cemented by EPS, are the frequent result. These deposits have significant effects on mutually cooperative growth of other bacteria and the formation and growth of pitting corrosion.9 The formation of Fe(OH)3 is accompanied by acidification (Figures 2 and 3) and increased corrosion, especially beneath the deposits. Certain neutrophiles, especially Gallionella, are wellknown for oxidizing Fe2+ autrophically at neutral pH. The versatility of thiobacilli, particularly T. ferrooxidans, is demonstrated in Figure 4a by the wide range of potentials and pHs where this remarkable species is active. T. ferrooxidans also acts as an anaerobic autotrophic sulfur oxidizer according to Table 1, which shows further that T. ferrooxidans and others are capable of autotrophic oxidization of Fe2+ to soluble Fe3+ at pH 3 or less. Fe3+ is well-known as a strong oxidizer and aggressive corrosion accelerator. Another strain, T. denitrificans, oxidizes sulfides and sulfur to sulfates while reducing nitrate to nitrogen anaerobically and autotrophically. The aerobic autotrophic oxidation of sulfide minerals to SO4- at low pH,4 also mediated by Thiobacillus sp., including the ever-effective T. ferrooxidans, is responsible for the notorious acid-waste problems that plague many operating and abandoned mine-waste streams throughout the world. Figure 4a shows that the area of thiobacterial (TB) activity extends to potentials sufficiently active to allow oxidation of reduced sulfur on

the Pourbaix diagram. The sulfide minerals are stable under anaerobic conditions within geologic mineral deposits but become susceptible to bio-oxidation when exposed to air in mines and waste piles. Coincident chemical or bio-oxidation of ferrous iron in the minerals is responsible for the formation of yellow Fe3+ water coloration and red Fe(OH)3 deposits in polluted stream beds. Alternatively, “metal-reducing” bacteria may anaerobically reduce and solubilize metal oxides when combined with more active half-cell reactions. For example, Table 1 shows that MnO2 may be reduced heterotrophically to soluble Mn2+ during simultaneous oxidation of organic matter (20), mediated by Bacillus sp. and others. Similarly, Fe(OH)3-based precipitates can be dissolved by heterotrophic reduction (7) to soluble Fe2+, again mediated by Bacillus sp. Autotrophic reduction of Fe(OH)3 (7) is also shown in Table 1 during oxidation of S0 to sulfate (14) and H2 to H+ (1). Shewanella putrafaciens has been demonstrated10 to effectively reduce Fe(OH)3 and MnO2, while oxidizing H2, as well as various organic compounds. Fe(OH)3 here represents the equilibrium form of various ferric oxide polymorphs, such as hematite, ferrihydrite, and goethite. Reduction of these ferric oxides has been associated with the removal of protective surface films and accelerated MIC.10-12 Sulfate-reducing bacteria (SRB) are well-known agents in the MIC of carbon and stainless steels as well as many nonferrous alloys. The area of SRB activity on the Pourbaix diagram (Figure 4b) coincides with the sulfur reactions from Figure 3. Certain SRB, especially Desulfovibrio sp., mediate the anaerobic reduction of SO42to S0 and/or H2S over wide ranges of pH, pressure, temperature, and salinity values4 by ((14)-(17)); they are well-known for aggressively accelerating corrosion. The mechanism of increased corrosion is still uncertain but probably involves both the enhanced formation of damaging H2S and the surface precipitation of catalytic surface films of iron-sulfide and elemental sulfur.13-15 Figures 3 and 4b show a narrow potential region where sulfur, S0, is stable between the reduction of SO42- to S0 by (15) and reduction of S0 to H2S by (17). S0 is wellknown for being especially corrosive to carbon steel and stainless steel.16 Newman et al.14,15 have suggested that S0, stable on the surface only in the narrow potential band shown in Figure 3, will catalyze anodic dissolution and hinder passivation within an acid pit or crevice beneath oxide or sulfide surface deposits. Furthermore, they have demonstrated that an external oxygen reduction cathode on neighboring oxide or sulfide surfaces drives anodic dissolution within the pit or crevice through a micro galvanic local occluded cell. Ironically, it would seem that nominally anaerobic SRB thrive best in at least partially aerobic conditions. Nearby, aerobic conditions, fostering cathodes near the anaerobic occluded anode, may be a necessity. It is notable that SRB are apparently inactive in very low pH solutions, according to Figure 1, supporting the observation that SRB have not been isolated from low-pH occluded cells with active MIC.23 Because the potential range for the sulfur reactions ((14)-(17)) is narrow, many SRB can mediate many if not all of the reduction reactions involving sulfur. For the same reason, sulfur and its anions can be readily oxidized aerobically ((14)-(17)) by reduction of dissolved oxygen (2), and anaerobically in the same half-cell reactions ((14)-(17)) by reduction of nitrate to nitrogen

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(10). Some references8 simply refer to sulfur-oxidizing bacteria, assuming that all such bacteria oxidize all reduced forms of sulfur to SO4-. It has been noted earlier that Thiobacillus sp. are well-known for their versatility to mediate both the aerobic and anaerobic oxidation of sulfur and its anions as well as Fe2+. It should be noted, as well, that the Pourbaix diagram is limited to reactions of equilibrium species. There have been recent observations17,18 of increased iron corrosion associated with microbially mediated reduction of thiosulfate, S2O32-, an important but metastable sulfur anion, which does not appear in Figure 3. It is likely, however, that the reduction of such species occur at redox potentials near those of their similar equilibrium counterparts. Figure 4c compares the various nitrogen reactions from Figure 3 with the activity areas for nitrogen reactions on the Pourbaix diagram. Nitrogen bacteria alternately oxidize and reduce nitrogen between the soil-water and atmosphere. Anaerobic dissimilatory nitrate reduction to nitrite, NO2-, (12), is the first step in denitrification in which nitrite is eventually reduced to nitrogen. Nitrite may be reduced subsequently to NH4+ as in (13). Nitrification reverses the reduction of nitrogen in the cycle where reduced nitrogen compounds are oxidized aerobically and autotrophically back to nitrite and nitrate in (13) and (12). Rhizobium sp., classified as heterotrophic anaerobic bacteria (HAB) in Figures 1 and 4c, directly convert (fix) atmospheric molecular N2 to NH4+ (18) under reducing conditions in biodegraded humus and within the roots of legumes. The resulting dissolved ammonium makes nitrogen, an essential nutrient, available for plant growth. NH4+ can also be formed in large quantities by the removal of amino groups for amino acids during the decay of organic (humic) substances. Copious NH4+, resulting especially from humic decay, complexes copper and causes stress corrosion cracking of Cu-Zn alloys (brasses), first reported in cartridge cases exposed to nearby rotting vegetation in the 1920s.19,20 Some HAB are acid-producing bacteria (APB) which generate corrosive acid metabolites. An anaerobic culture of mixed Clostridium sp. caused carbon-steel corrosion which was simulated in subsequent abiotic tests by exposure to the acetic and lactic acids analyzed in the metabolite solutions.21 Daniels et al.22,23 reported increased anaerobic corrosion of iron in the presence of several methanogens which form methane from carbon dioxide by (20) and fall also in the activity area for HAB, as shown in Figure 4d. They theorized that the methanogens increased corrosion in their experiments by reacting with and removing surface hydrogen which acts as a reaction barrier. The methanogens thereby “depolarize” the hydrogen reaction and increase the corrosion rate. They claimed confirmation when CH4 could be measured in a vessel in which methanogens reacted with H2 supplied from corroding Fe in an adjoining abiotic vessel.23 However, these dual-vessel experiments show only that CH4 can form from corrosion-generated H2. They do not rule out the possibility of an alternative, that is, the direct increase of iron oxidation (3) by biomediated cathodic reduction of CO2 (20) to form CH4. Hydrogendepolarization theory has been largely discredited in the corrosion literature. Instead, it has been demonstrated that an oxidizer with a more noble half-cell electrode potential such as (19) acts directly to increase the anodic

half-cell dissolution reaction (3), decreasing hydrogen evolution without any necessary direct reaction between hydrogen and the oxidizer.24,25 In any event, any of the bacteria listed in Table 1, which mediate anaerobic hydrogen oxidation (1), might also substitute iron anodic dissolution (3), nearby in Figure 4d, to directly affect the corrosion of iron. The previous discussion clarifies Figure 1, which defines areas of potential and pH where various bacteria are active. The iron bacteria (IB) in Figure 4a autotrophically oxidize ferrous oxide or soluble Fe2+ in aerated neutral or acidic solutions to insoluble ferric oxides by (6) or (7), respectively. The highly versatile thiobacteria (TB) autotrophically oxidize ferrous iron and sulfur compounds, mostly in aerated acidic solutions (Figure 4b). Finally, the SRB, especially important in MIC, are a subset of the HAB, which specialize in the reduction of sulfate to form H2S in neutral to slightly acidic solutions (Figure 4d). The HAB in Figure 4c include those discussed here which reduce sulfur compounds and/or carbon dioxide while oxidizing organic carbon in anaerobic solutions. Figure 4c compares the various nitrogen reactions from Figure 3 with the activity areas for nitrogen reactions on the Pourbaix diagram. Nitrogen bacteria alternately oxidize and reduce nitrogen, the reactions of which form the nitrogen cycle. These reactions take place in soil, sediment, and aquatic environments. Anaerobic dissimilatory nitrate reduction to nitrite, NO2- (12), is the first step in denitrification in which nitrite is subsequently reduced to nitrogen (10). Cells may also reduce nitrite in an assimilatory manner for biosynthesis of cellular components: proteins and nucleic acids (13). Nitrification reverses the reduction of nitrogen in the cycle, where reduced nitrogen compounds are oxidized aerobically and autotrophically back to nitrite (12) and nitrate (13). No apparent biological activity involving nitrogen is apparent in Figure 4c below pH 6. Below pH 3.3, the lines for (12) and (13) cross over one another and approximately represent half-cell reactions involving nitrous acid, HNO2.26 Summary and Conclusions The charge-transfer (electrochemical) equilibria of microbiological cell growth and iron corrosion have been instructively compared on the potential-pH (Pourbaix) diagram. Microorganisms which mediate numerous electrochemical cell reactions have been tabulated and discussed relative to their significance in microbial ecology and microbiologically influenced corrosion. The discussion highlights the diversity of available bacteria which mediate electrochemical half-cell reactions over the wide range of redox conditions from oxidizing (aerated) to reducing (deaerated) and acidity from pH 1.5 to 11, with or without oxidation of organic nutrients. Any given half-cell reaction can either be oxidized or reduced by specific but different bacteria, depending on the redox conditions present. The comparisons confirm experimental observations that corrosion can be enhanced by sulfate reduction to form H2S and iron sulfides, solubilization of surface-oxide passive films, and formation of corrosive acids. All may be further enhanced by aerobic biomediated formation of biofilms and insoluble surface deposits which form occluded acidic anaerobic micro environments and concentration cells. Iron corrosion has been increased also by methane-

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forming bacteria (methanogens), and the present analysis suggests that the mechanism may involve direct anodic oxidation of Fe to Fe2+ combined with cathodic reduction of CO2 to CH4, without intermediate oxidation of H2 (hydrogen depolarization). It seems possible that other hydrogen-oxidizing bacteria besides methanogens may directly influence corrosion by substituting Fe anodic dissolution for hydrogen oxidation. Acknowledgment This work was sponsored by the Lawrence Livermore National Laboratory by a contract to the University and Community College System of Nevada. We are grateful to W. L. Clarke, Jr., R. D. McCright, and D. Stahl for continuing support and interest. Literature Cited (1) Baas Becking, I. G. M.; Kaplan, I. R.; Moore, D. J. Geol. 1960, 68, 234. (2) Pourbaix, M. Atlas of Electrochemical Equilibria in Aqueous Solutions; NACE: Houston, TX, 1974; p 307. (3) Atlas, R. M.; Bartha, R. Microbial Ecology, 3rd ed.; Benjamin Cummings: New York, 1993: p 328. (4) Ehrlich, H. L. Geomicrobiology, 3rd ed.; Marcel Dekker: New York, 1996. (5) Geesey, G. In A Review of the Potential for Microbially Influenced Corrosion of High-Level Nuclear Waste Containers; Cragnolino, G. A., Ed.; CNWRA 93-014, Nuclear Regulatory Commission Contract NRC-02-88-005; U.S. Government Printing Office: Washington, DC, 1993. (6) Lindsay, W. L. Chemical Equilibria in Soils; Wiley: New York, 1979; p 374. (7) Lehninger, A. L. Biochemistry, 2nd ed.; Worth, New York, 1975; pp 443, 479. (8) Laskin, A. I.; Lechevelier, H. A.; Johnson, C. L. Handbook of Microbiology, 2nd ed.; CRC Press: Cleveland, 1977: Vol. 1, pp 237-245, 285-307. (9) Dexter, S. C. Metals Handbook, Vol. 13, Corrosion, 9th ed.; ASM International: Metals Park, OH, 1987; p 114.

(10) Nealson, K. H.; Saffarini, D. A. Annu. Rev. Microbiol. 1994, 48, 311. (11) Obuekwe, C. O.; Westlake, D. W. S.; Plambeck, J. A.; Cook, F. D. Corrosion 1981, 37, 461. (12) Little, B.; Wagner, P.; Hart, K.; Ray, R.; Lavoie; Nealson, K.; Aguilar, C. CORROSION/97; NACE: Houston, TX, 1997; Paper No. 215. (13) Lee, W.; Lewandowski, Z.; Nielsen, P. H.; Hamilton, W. A. Biofouling 1995, 8, 165. (14) Webster, B. J.; Newman, R. C. Corros. Sci. 1993, 35, 675. (15) Newman, R. C.; Webster, B. J.; Kelly, O. J. ISIJ Int. 1991, 2, 201. (16) Schmitt, G. Corrosion 1991, 47, 285. (17) Margot, M.; Carreau, L.; Cayol, J.-L.; Olivier, B.; Crolet, J.-L. In Microbial Corrosion, Proceedings 3rd European Federation Corrosion Workshop, Portugal, 1994; Institute of Materials: London, 1995; p 293. (18) Crolet, J.-L.; Margot, M. CORROSION/95; NACE: Houston, TX, 1995; Paper No. 188. (19) Shreir, L. L. Corrosion, 2nd ed.; Newnes-Butterworths: London, 1976; Vol. 1, 4:56. (20) Jones, D. A. Principles and Prevention of Corrosion, 2nd ed.; Prentice-Hall: Englewood Cliffs, NJ, 1996; p 236. (21) Pope, D. H.; Zintel, T. P.; Kuruvilla, A. K.; Siebert, O. W. CORROSION/88; NACE: Houston, TX, 1988; Paper No. 79. (22) Boopathy, R.; Daniels, L. Appl. Environ. Microbiol. 1991, 57, 2104. (23) Daniels, L.; Belay, N.; Rajagopal, B. S.; Weimer, P. J. Science 1987, 237, 509. (24) Fontana, M. G.; Greene, N. D. Corrosion Engineering, 2nd ed.; McGraw-Hill: New York, 1986; p 316. (25) Jones, D. A. Principles and Prevention of Corrosion, 2nd ed., Prentice-Hall: Englewood Cliffs, NJ, 1996: p 90. (26) Pourbaix, M. Atlas of Electrochemical Equilibria in Aqueous Solutions; NACE: Houston, TX, 1974; p 499.

Received for review July 14, 1999 Revised manuscript received February 1, 2000 Accepted February 4, 2000 IE990516V