In addition to its metallic elemental form (Fe0), iron has only the +2 ferrous, Fe(II), and +3 ferric, Fe(III), oxidation states, separated by a one-electron difference. Most iron cycling involves oxidations and reductions between these two states, via both abiotic and biochemical mechanisms.
Iron is one of the most abundant elements on Earth, but it is confined primarily to the lithosphere (soil averages about 7% Fe). It is nonvolatile, and its concentration in air is virtually zero except for the small amounts present in airborne soil particles. In water, the solubility of ferric iron is low (<0.02 mg/L at pH 6), except under very acidic conditions or in some organic complexes. Ferrous iron is more soluble (about 20 mg/L at pH 6, but depending strongly on alkalinity). However, under the anaerobic conditions in which it is usually found, it is likely to form a precipitate with sulfide. Although iron is an essential element, it typically is present in organisms only at low concentrations (e.g., 0.2% of the dry weight of E. coli, Table 11.1). Thus, although iron is abundant and only small amounts are needed, it still may be available only at levels that limit growth in some environments, especially the open ocean.
In the absence of oxygen and nitrate, ferric iron can be used for respiration by a variety of aerobic bacteria. In the process, the Fe(III) is reduced to Fe(II). Both organotrophic and lithotrophic iron reducers are known. The process seems to be carried out by many of the same organisms that can reduce nitrate.
There are a number of natural oxygen-deficient conditions under which reduction might take place. These include anoxic biofilm layers, lake bottoms and sediments, water-saturated soils, groundwaters, and bogs. Reduction is a principal means of solubi-lizing iron-bearing solids. Iron reduction and resolubilization also occur in landfills and anaerobic digesters. Where sulfides are present, FeS will form, precipitating the iron while helping to limit odor and toxicity problems associated with H2S.
A diverse assortment of aerobic bacteria and archaea can oxidize ferrous iron, apparently using it as an energy source. Because only a small amount of energy is available from this transformation, large amounts of iron must be processed to support their growth.
At neutral pH values, ferrous iron is unstable under aerobic conditions and will be oxidized rapidly and spontaneously (abiotically) to the +3 form. Thus, microorganisms utilizing iron at such pH values must be located at the interface between aerobic and anaerobic zones, depending on this specialized niche so that they may capture this energy before it is lost. On the other hand, at low pH, ferrous iron is sufficiently stable that microorganisms can utilize it in aerobic systems.
Thus, iron oxidizers can be subdivided into two groups based on their pH preferences. The acidophiles include several Thiobacillus, Leptospirillum (a xenobacteria), and Sulfo-lobus (an archaea). Many of the same organisms can also oxidize reduced sulfur, an advantage since ferrous iron and sulfides often occur together. T ferrooxidans, an autotroph, is the best known; it has a pH range of 0.5 to 6, with an optimum at 2! It utilizes a copper-containing protein, rusticyanin, for transferring the electron from iron during its oxidation (Figure 13.27). The amount of energy available from the oxidation is too small to use directly for ATP synthesis; instead, the proton gradient is utilized for this purpose.
Gallionella (Figure 13.28) grows at neutral pH, usually at a point where anaerobic groundwater comes in contact with oxygen (such as in wells). It is commonly heavily coated with ferric hydroxide, Fe(OH)3. Sphaerotilus and Leptothrix also deposit oxidized iron on their sheath, but it is not certain that they are able to capture the energy from iron-oxidation. Other iron, oxidizing bacteria include Planctomyces and Hyphomicrobium;
Elevated at low pH conditions
Cytochrome c Cytochrome a1
Cytochrome c Cytochrome a1
Cytoplasmic H+ sink (more neutral pH)
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