Iron and Trace Element Chemistry: Fe2+/Fe3+, Chelators and Micronutrients

Why iron precipitates as pH and oxygen rise, how fast ferrous iron oxidises, how EDTA and DTPA chelates hold iron at different pH and light levels, and what each trace element does.

Ferrous and ferric iron

Iron in natural water occurs almost entirely in two oxidation states. Ferrous iron, Fe2+, is the reduced form, and ferric iron, Fe3+, is the oxidised form. Iron is the fourth most abundant element in the Earth's crust, yet it is scarce in oxygenated surface water. Fe2+ reacts readily with oxygen, and Fe3+ has very low solubility, so iron limits biological productivity in many waters. Dissolved ferrous iron is invisible. Once oxidised, the iron precipitates and appears as suspended reddish-brown particles.

Boyd notes that in oxygenated water, pH controls how much iron stays dissolved. Ferric iron seldom exceeds 2 mg/L as Fe unless the pH is below 4. Fresh waters can still carry up to about 1 mg/L or more of dissolved iron, because iron forms soluble hydroxide species, ion pairs and complexes with dissolved organic matter. Iron(III) oxyhydroxide precipitates from ferric salt solutions between pH 6.5 and 8. Under natural conditions, Fe2+ is about 100 times more soluble than Fe3+ and is the form most organisms take up most easily.

How fast ferrous iron oxidises

The conversion Fe2+ → Fe3+ → Fe(OH)3 depends mainly on pH, temperature and dissolved oxygen, and also on other ions in solution. Rutgers Cooperative Extension gives the following times for 90% of ferrous iron to oxidise. At pH 7.0 it takes about 1 hour at 21 °C and about 10 hours at 5 °C. At pH 8.0 it takes about 30 seconds, and at pH 6.0 about 100 hours. Below about 2 mg/L dissolved oxygen, oxidation is very slow. UF/IFAS Extension states that ferrous iron is readily oxidised to the ferric form once pH rises above 5.3. When dissolved ferrous bicarbonate oxidises, it forms ferric hydroxide, releases carbon dioxide and lowers the pH.

These rates explain several common observations. Groundwater with no dissolved oxygen can carry 20 mg/L or more of ferrous iron. When such water is aerated, the iron drops out as a hydroxide floc. In ponds the floc mostly settles harmlessly, but in hatcheries and intensive tanks it can coat eggs and clog the gills of fish and shrimp. Iron-oxidising bacteria, mainly filamentous genera such as Gallionella, Leptothrix and Sphaerotilus, oxidise Fe2+ and form slimy deposits. Ferric hydroxides also bind phosphate tightly, which removes it from the water. Iron oxide-hydroxide is used in aquarium water treatment as a phosphate binder for this reason. Under anoxic conditions, microbes reduce Fe(III) back to Fe2+ and the bound phosphorus is released again.

Chelators: EDTA and DTPA

A chelator is a ligand that binds a single metal atom through two or more bonds and so forms a ring around it. Chelating ligands bind a metal more strongly than comparable single-bond ligands; this is called the chelate effect, and it arises mainly from entropy. In nature, most dissolved metal complexes are chelates with humic substances, proteins or microbial siderophores. Grasses short of iron release their own chelator, the amino acid mugineic acid, from their roots.

EDTA (ethylenediaminetetraacetic acid) binds both iron and calcium and forms water-soluble complexes even at neutral pH. The EDTA anion wraps around the metal through four carboxylate groups and two amine groups. Near neutral pH the main iron complex is [Fe(EDTA)(H2O)]-, which gives a yellow-brown solution, while the reduced Fe(II)-EDTA is almost colourless and oxidises very easily. Without EDTA or a similar chelator, ferric ions form insoluble solids that plants cannot use. DTPA (diethylenetriaminepentaacetic acid) has five carboxymethyl groups on a diethylenetriamine backbone and acts as an expanded version of EDTA. Iron-DTPA is used as an aquarium plant fertiliser because it keeps iron from precipitating as Fe(OH)3. It is unclear how far DTPA protects dissolved Fe2+ from oxidation by air.

How well a chelate holds on depends on pH. In a UF/IFAS Extension comparison, all three common iron chelates are fully chelated at pH 6.0. At pH 7.5, the fraction of iron still chelated is 1.0 for EDDHA, 0.5 for DTPA and only 0.025 for EDTA. Older hydroponic work found no iron loss from Fe-EDTA in a nutrient solution below pH 6. EDTA therefore holds iron reliably only in slightly acidic water, while DTPA keeps it chelated to a higher pH.

Light also breaks down iron chelates. Iron(III)-EDTA is removed from surface waters mainly by direct photolysis at wavelengths below 400 nm, with half-lives from about 11 minutes to more than 100 hours depending on the light. EDDHA is applied in the evening in agriculture to avoid this photodegradation. Plants take up the iron and the ligand separately: the stable ferric chelate is first reduced to the less stable ferrous chelate at the root surface.

Trace elements and their roles

UF/IFAS Extension defines micronutrients as elements present in plant tissue at 100 ppm or less: iron, zinc, manganese, copper, boron, chlorine, molybdenum and nickel. Of these, iron, manganese, zinc and copper are the ones that oxidise or precipitate easily. Their availability falls once pH rises above about 6.5, which is why they are the ones usually supplied as chelates.

Iron: not part of the chlorophyll molecule but essential for making it; also needed for electron transport and as an enzyme cofactor. Iron does not move within the plant, so deficiency shows first as yellowing between the veins of new leaves.
Manganese: needed for photosynthesis, including chloroplast formation. An excess of manganese can interfere with iron uptake.
Zinc: required by many enzymes and has an essential role in DNA transcription.
Copper: involved in photosynthesis, many enzymes and lignin formation. Copper deficiency can promote iron deficiency.
Molybdenum: part of nitrate reductase and nitrogenase, so it is central to nitrogen metabolism.
Nickel: activates urease, the enzyme that processes urea.
Boron: needed for cell-wall formation and for moving sugars across membranes.
Chlorine (as chloride): needed for osmotic and ionic balance and has a role in photosynthesis.

Deficiency symptoms are covered in the plant nutrient deficiency guide, and metal toxicity to fish and invertebrates is covered in the heavy metals guide. In pond aquaculture, iron itself is almost never present at concentrations high enough to poison fish or shrimp directly. The problems it causes are physical, such as floc on eggs and gills, and indirect, such as oxygen consumption when ferrous iron oxidises after a stratified pond mixes.