



Biogeochemical cycles describe pathways by which chemical elements move through both biotic (the biosphere) and abiotic compartments (the atmosphere, hydrosphere, and lithosphere) on Earth. Along with energy flows, biogeochemical cycles establish the relations among ecosystem compartments at local, regional and global scales. In these systems of inputs, outputs, sources and sinks, elements are moved from one part of an ecosystem (e.g., ocean, soil, atmosphere) where the element may temporarily accumulate to another, back and forth among organisms, and from living organisms to the abiotic environment and back again. In other words, chemical elements are cycled and reused within and among Earth's various compartments over and over again.

The biogeochemical cycles proceed through biological, geological and chemical interactions along hydrological, gaseous, and mineral "trade routes." Among the most ecologically important and well known are the element cycles of carbon (C), nitrogen (N), oxygen (O), phosphorus (P), and sulfur (S), as well as the water (H 2 O) cycle. One biogeochemical cycle that is often overlooked, however, is Earth's iron (Fe) cycle (Figure 1).



Figure 1: The biogeochemical cycle of iron on Earth Natural processes, anthropogenic activities, and microbial communities affect the iron cycle. In nature, both iron oxidative and reductive reactions (a) depend on environmental conditions and microbial activities. Iron is oxidized into Fe3+ under aerobic conditions, or by microbes under acidic pH, and it is reduced to Fe2+ under anaerobic conditions. Iron can be transported to the ocean (b) as dust or volcanic ash. Coarse particles will sink rapidly, while smaller (colloidal) particles will travel further and stay in the surface ocean, increasing the amount of bioavailable iron (e.g., Duggen et al. 2009). Siderophores (c) are chelating agents secreted by microorganisms under low iron stress. They scavenge Fe from the environment and form Fe(III)-siderophore complexes, making the otherwise insoluble iron, available for bacterial uptake. The complex can be photolysed, resulting in the reduction of Fe(II) and a ligand, thus becoming available for planktonic communities (e.g., Barbeau et al. 2001). © 2010 All rights reserved.

Iron composes more than 30% of the Earth's mass, and is a ubiquitous element found in the atmosphere, biosphere, lithosphere, and hydrosphere. It is one of the most abundant elements on Earth and among the most important elements in the biosphere (Morgan & Anders 1980). It is an essential element for countless cellular processes and metabolic pathways in both eukaryotic and prokaryotic organisms. Yet despite its abundance, iron can be in short supply for growing organisms because it changes its chemical form, in ways that govern its availability. In its pure state it is a reactive metal that oxidizes readily in the presence of oxygen. On Earth, iron exists in one of two oxidation states: reduced ferrous iron, depicted as Fe(II) or Fe2+, or oxidized ferric iron, depicted as Fe(III) or Fe3+). These states can be found in nature as solids in the form of Fe(III)- and Fe(II)-bearing minerals (Figure 2) or as ions (Fe3+ and Fe2+) dissolved in water. The amount of Fe in a body of water, and the prevalence of each oxidation state, are controlled by oxygen concentration, pH, and the biological activities of microorganisms and higher organisms. The biogeochemistry of exchange between iron's two forms has been amusingly referred to as Earth's "ferrous wheel."



Figure 2: Iron-containing minerals Iron oxides such as hematite (a and d) and goethite (b and e) are among the most common minerals found on Earth. Pyrite (c and f) is an iron sulfide known as "fool's gold" because of its resemblance to gold. When exposed to the atmosphere during mining, pyrite react with oxygen and water to form sulfate (SO 4 2-) and H+ ions resulting in acid mine drainage (AMD). © 2010 All rights reserved.

On Earth when water comes into contact with solid iron in the presence of atmospheric oxygen, Fe is rapidly oxidized to Fe2+ via the reaction:

Fe (solid) → Fe2+ (aqueous) + 2e-

Hydroxide ions in the water then react with the Fe2+ ions to produce solid Fe(III)-minerals through the reaction:

Fe2+ (aqueous) + 2OH- (aqueous) → Fe(OH) 2 (solid mineral made of ferric iron)

These Fe(III)-minerals are often powdery and red in color when dried, and are commonly referred to as rust. Different Fe(III)-minerals (Table 1 & Figure 2) will form under different environmental conditions (e.g., oxygen exposure, pH, presence of other ions). Oxidized, particulate forms of iron are poor sources for the growth of aquatic life because iron, in its solid form, is not readily available for organisms to use for growth and reproduction.

Table 1. Common iron-containing minerals found on Earth Oxides Oxy-hyrdoxides Hydroxides Hematite α- Fe 2 O 3 Akaganéite β-FeOOH Bernalite Fe(OH) 3 Maghemite β-Fe 2 O 3 Feroxyhite δ'-FeOOH Fe(OH) 2 γ-Fe 2 O 3 Ferrihydrite Fe 5 HO 8 + 4H 2 O Green rust1 ε-Fe 2 O 3 Goethite α-FeOOH Magnetite Fe 3 O 4 Lepidocrocite γ-FeOOH Wüstite FeO Schwertmannite δ-FeOOH 1. In general, green rusts have the formula Fe3+ x Fe2+ y (OH) 3x+2y-z (A-); where A- = Cl- or ½ SO 4 2-

Numerous iron-containing minerals exist on Earth (Table 1 & Figure 2). These minerals are classified as Fe(III)-oxides (e.g., hematite), Fe(III)-oxyhydroxides and -hydroxides (e.g., goethite), Fe(II)-minerals (e.g., pyrite), or mixed Fe(II)-Fe(III) minerals (e.g., magnetite). The formation of these inorganic compounds initially involves the aerobic weathering of primary rocks in terrestrial and marine systems (Cornell & Schwertmann 2003). This process might be followed by redistribution between compartments (e.g., hydrosphere and lithosphere) and may involve mechanical transport by wind and/or water (Cornell & Schwertmann 2003). Iron-metabolizing microorganisms, which catalyze the oxidation or reduction of iron, also play a significant role in the cycling of iron in the environment and between compartments.

Fe(III)-minerals are characterized by low solubility at circumneutral pH, and thus exist as a solid in much of Earth's near-surface environments. However, in environments that are strongly alkaline or acidic, Fe(III)-minerals can dissolve because of the mineral's ability to act both as a base and as an acid. Fe(II)-minerals, on the other hand, are considerably more soluble at neutral pH, and thus readily dissolve to release Fe2+ ions to the surrounding environment. Fe(II)-minerals are only stable at neutral or alkaline pH in anoxic environments (i.e., when no oxygen is present). In an aqueous or moist environment, exposed to oxygen from Earth's atmosphere, Fe(II) rapidly oxidizes to Fe(III), with a half-life of several minutes (Stumm & Morgan 1996). The Fe(III) that is produced exists as an Fe3+ ion in acidic environments, or as solid Fe(III)-minerals (Figure 2) in neutral or slightly alkaline environments.

The movement of iron through the biosphere is controlled, at least in part, by the actions of plants. Although considered a micronutrient for plants, iron is an essential trace element required for the production of chlorophyll. Since iron uptake occurs at the tip of the roots, anything that interferes with this process will result in iron deficiency. An iron deficient plant will show a yellowish color in younger leaves. However, while small amounts are necessary for growth, iron can become toxic to plants. Iron toxicity is associated with large concentrations of Fe2+ in the soil solution (Becker & Asch 2005), and has been described as one of the major growth limiting factors for rice (Dobermann & Fairhurst 2000, Sahrawat 1979, Sahrawat 2004). Iron toxicity occurs in inundated soils with high iron-oxidizing activity, although the accumulation of organic acids, hydrogen sulfide and other reduction products might contribute to this toxicity (Sahrawat 2004).

Because iron is an important micronutrient used by most organisms, and is required for important cellular processes such as respiration, oxygen transport in the blood, photosynthesis, nitrogen fixation, and nitrate reduction, its bioavailability is of concern for Earth's living organisms, especially in aquatic ecosystems. This is because, despite its relatively high abundance on Earth, iron is a minor component of aquatic systems because of its relative insolubility in water at circumneutral pH. Research has demonstrated that most areas of the open ocean have surface trace metal concentrations in the nanomolar to picomolar range (Morel & Price 2003). That means one iron ion present for every fifty billion or trillion water molecules. At that vanishingly low concentration, iron is very often the limiting nutrient for primary production (e.g., photosynthesis) in large expanses of ocean (Behrenfeld & Kolber 1999, Coale et al. 1996, Hutchins & Bruland 1998, Martin et al. 1994).

In the upper ocean, dissolved iron frequently occurs in the form of complexes with strong ligands (Rue & Bruland 1995) presumably of biological origin (Hutchins et al. 1999, Kondo et al. 2008, Mawji et al. 2008). A ligand is an ion or molecule that binds to a metal, such as iron. A study by Barbeau et al. (2001), showed that Fe(III)-binding molecules synthesized by ocean-inhabiting microorganisms might facilitate the photochemical cycling of iron in ocean surface waters. Furthermore, their research demonstrated that key factors of the cycle included photolysis of the Fe3+-ligand complexes (e.g., siderophores, which are metal chelating molecules synthesized by some microorganisms), the reduction of Fe3+ to Fe2+, and oxidation of the ligand, thus increasing the bioavailability of iron for uptake by planktonic communities (Figure 1) (Barbeau et al. 2001). Therefore, the presence of microbial populations is key to an understanding of iron cycling in the ocean, because they produce the organic ligands that can strongly affect the solubility of iron, and thus the iron uptake mechanisms utilized by numerous marine organisms (Geider 1999).