Registration Dossier

Environmental fate & pathways

Endpoint summary

Administrative data

Description of key information

Additional information

In general, degradation is an irrelevant process for inorganic substances that are assessed on an elemental basis.Under REACH (ECHA 2008, Chapter R.7B – Endpoint Specific Guidance), the term ‘Hydrolysis’ refers to the “Decomposition or degradation of a chemical by reaction with water”, and this is a function of pH (i.e. abiotic degradation) for the respective metal kations present in the submission item.

The behaviour of ions of iron (Fe2+ and Fe3+) in solution is dependent on the different conditions. Key conditions that influence iron behaviours are the oxygen content, pH, presence of potential ligand anions with which the cationic iron might associate. Iron exists in numerous forms in water and is commonly bonded or co-ordinated to other species such as water molecules or electron donor partners (Johnson et al 2007).


Iron in the ferric form is subject to phototransformations. Photoreduction of ferric (Fe3+) ions to ferrous (Fe2+) ions is evidenced in fresh and marine waters and is subject to a number of publications (Barbeau 2006, Emmenegger et al 2001, Hong & Kester 1986, Kopácek et al 2005, Rose & Waite 2003). Ferrous sulfate and ferrous chloride are reducing agents, and exposure to light gives reduction of the ferric form. The process of photoreduction may overcompensate the aerobic oxidation resulting in the presence of ferrous (Fe2+) ions during daylight and disappearing during the night (Kopácek et al 2005). Accordingly low concentrations of the ferrous kation apart from the top metre of water where light can penetrate and apart from the very deep anaerobic environments can be expected as shown by Hong & Kester (1982).

Photoliberation (photodegradation of organically bound iron (III) complexes typical to boreal forests and tundra soil) of organically bound iron under the influence of 300 nm light resulting in kationic forms with a half-life of less than 12 h was shown in laboratory experiments (Kopácek et al 2005).


Both ferrous and ferric ions released into (or generated in) water will rapidly precipitate as highly insoluble oxides and oxo-hydroxides. Any precipitate would in turn undergo further oxidation to form the highly insoluble ferric hydroxide (Fe(OH)3) as this species autocatylzes the oxidation of Fe(II). Thus the formation of ferric hydroxide at pH levels above 5.0 limits the presence of iron in aqueous systems. The resulting ferric hydroxide is very insoluble and largely confined to particulate phases.

These stable compounds are exactly the forms in which iron is found naturally in the earth’s crust. Above pH 2, colloidal gels are formed, giving a precipitate of the red–brown gelatinous hydrous iron oxide. These insoluble species are thus predominant at environmentally relevant pH. With time, these hydroxides either polymerise to form larger insoluble stable complexes or they are trapped and buried in sediments (IHCP 2009, p 511 & ECHA 2011, p 475).

Reactions with oxygen

In water containing oxygen, the chemistry is dominated by iron-oxygen reactions. These are highly complex mechanistically (Cotton & Wilkinson 1972), however, the end points are clear. Ferrous ion, Fe(II), is unstable when its solutions are exposed to air, and it oxidizes to the ferric ion, Fe(III), which then forms the familiar insoluble, hydrated, amorphous, gelatinous precipitate, Fe(OH)3 (ferric hydroxide). Other common oxidant such as nitrate can oxidize iron(II) as well (Schnitzer 1969). The rate at which these conversions takes place is important because if Fe(II) and Fe(III) ions are rapidly removed from solution as insoluble precipitates, then any direct impact of dissolved Fe ions on the aquatic environment will be minimal. Ferrous iron is also a powerful catalytic species, which causes the decomposition of organic matter in the presence of oxygen. With some substrates, only traces of Fe(II) may cause the catalytic effect (Kopácek et al 2005).

Ferrous ions in aqueous solution readily oxidize to the ferric form according to the rate equation:

-d[FeII]/dt = k[FeII][O2][OH-]2

Where k = 1.5 x 10^16 L^-3 mole^-3 min^-1,

[FeII] = concentration of ferrous ion in solution,

[O2] = concentration of oxygen in solution,

[OH-] = hydroxyl ion activity.

Thus for an iron salt at a concentration of 10 mg/L Fe at pH 7.0 and saturated oxygen levels an initial rapid conversion rate of 3.6 x 10^-4 mol L^-1 min^-1 (20 mg L-1 min-1) is calculated (Skeaff 2004).

The rates at which dissolved ferrous sulphate (Fe2+) oxidizes to (Fe3+) and forms the ferric hydroxide [Fe(OH)3] precipitate (Skeaff 2004):

  • is highly dependent on pH (100 fold increase with a unit from pH 6 to 8)
  • decreases with an increase in ionic strength of the aqueous medium (pristine waters will contain less iron)
  • is dependent to some extent on the anions present in solution such as sulphate and chloride
  • increases 10-fold for a 15 °C increase in temperature
  • exhibits a linear dependence on the partial pressure of oxygen, and
  • is dependent on the initial concentration of ferrous sulphate and exhibits linear reaction kinetics at Fe(II) loadings less than ~50 micromolar (~3 mg/L). At concentrations greater than 50 micromolar, rates of reaction increase with increasing concentration of ferrous sulphate (about 4x for each order of magnitude).

Implications for understanding behaviour in water

Based on literature data and empirical reaction kinetics, it can be calculated that, at pH 6 in the OECD TG 203 medium (diluted by 10 as per the OECD Transformation/Dissolution Protocol), the half-times for the oxidation of Fe(II) are 11, 9 and 3.6 h, for 1, 10 and 100 mg/L loadings of FeSO4, respectively. At pH 8, the reaction is estimated to be as short as 8 seconds (Skeaff 2004).

Table: Half-times for the oxidation at pH 6 of Fe(II) as FeSO4

Loading [mg/L]

Half-time [h]

 1 11
 10 9.0
 100 3.6

For substances where the dissipation half-life (DT50) is less than 12 hours, environmental effects are likely to be attributed to the transformation products rather than to the parent substance itself (ECHA 2008 Information requirements and chemical safety assessment Chapter R.7b: Endpoint specific guidance, Section R7.8, p 190). Thus the Chemical Safety Assessment (CSA) bases on Fe(III) effects.

The rapid precipitation of iron from aqueous systems accounts for low iron concentrations found in most natural aquatic systems (all except natural waters at very low pH values (i.e. < pH 5.5). At pH 6 and a low initial concentration of 1 mg/L FeSO4, 70 % removal from solution is calculated to be achieved in 19 h and 90 % removal would be achieved by 36 h. This can be considered as rapid degradability as it is well within the 28 day time period specified in OECD guidance documents and indicates that there is no concern in terms of long term environmental effects.

In natural ecosystems the absence of oxygen or low pH can result in iron salts remaining in solution but under such conditions environmental effects would be strongly influenced by these parameters. The presence of other ions in solution, such as carbonates and humates, is expected to stabilize ferrous but this is not expected to be a sufficient effect to overcome the precipitation.

Natural Organic Matter (NOM)

The role of natural organic reductants in environmental ecosystems is difficult to characterise because most natural organic matter is of indeterminate composition. However, the possibility that high molecular weight Natural Organic Matter (NOM) acts as a reductant in environmental systems (particularly anoxic ones) is widely acknowledged (Tratnyek & Macalady 2000). It is believed that the reducing potential of NOM is due to specific moieties such as complexed metals or conjugated polyphenols. Often, redox reactions involving these moieties are reversible, which means that NOM often serves as a mediator of redox reactions rather than being just an electron donor or acceptor.

Like the various forms of iron, NOM apparently serves as both bulk reductant and mediator of reduction. NOM can also act as an electron acceptor for microbial respiration by iron-reducing bacteria, thereby facilitating the catabolism of aromatic hydrocarbons under anaerobic conditions. In general, it appears that NOM can mediate electron transfer between a wide range of donors and acceptors in environmental systems. In this way, NOM probably facilitates many redox reactions that are favourable in a thermodynamic sense but do not occur by direct interaction between donor and acceptor due to unfavourable kinetics (Tratnyek & Macalady 2000).

Significance of humic substances in soil, surface waters and sediment / bioavailability

The ubiquity of iron and humic substances in the environment necessitates some discussion of the interactions between these species. Humic substances are major constituents of soil organic matter humus that contributes to soil chemical and physical quality. They can also be found in peat, coal, many upland streams and ocean water. Some details of their chemical constitution have been reviewed (Schwarzenbach et al 1993). The term “humic substances” includes humic acid (base-extractable) and fulvic acid (acid- and base-extractable). Humin and kerogen are not base-extractable. They are all part of a heterogeneous supramolecular system of bio-organic molecules, each having molecular mass around 2000 Da.

In natural waters complexes with certain organic molecules greatly alter solubility and bioavailability of iron. Many organic acids form strong soluble complexes with ferrous and ferric ions. An enrichment of iron is commonly found in surface waters with a high content of dissolved organic matter. These high concentrations of complex soluble iron are associated with high levels of humic acids, tannic acids and other lignin derivatives (Wetzel 1983). Effects of such external iron loading from Northern Finnish Rivers can be seen in elevated iron and colour levels in the Oulu estuarine, especially during winter and spring flooding (Hilli & Pienimäki 2003; see considerations on environmental background concentrations in the see discussion of environmental fate and pathways). The extent of organic complexation of iron in sea surface water is estimated to be approx. 99 % complexed in oceanic/coastal sea water (Morel & Hering 1993).

A large number of humic molecules are represented by hydrophobic compounds (long alkyl-chain alkanes, alkenes, fatty acids, sterols, terpenoids, and phenyl-alkyl residues of lignin degradation), which allow their self-association into supramolecular structures separated from the water medium. Humic substances have acidic functional groups, mainly carboxylic acids, and also phenolics, which confer on these molecules the ability to chelate multivalent kations such as iron ions. Typical ratios of C:H:O are around 10:12:6 (Schwarzenbach et al 1993). This chelation of ions is an important role of humic acids with respect to living systems. The uptake of these ions is facilitated by the prevention of their precipitation, and increasing their bioavailability, although the high molecular weights prevent direct uptake of entire molecules.

In natural systems, bacteria interact with the humic redox system, further complicating the understanding of the chemical processes taking place. Microbial reduction of humic acids and subsequent chemical reduction of poorly soluble Fe(III) minerals by the reduced humic acids represents an important path of electron flow in anoxic natural environments such as freshwater sediments as well as soil (Kappler et al 2004).

The interaction between iron and humic substances is not straightforward. For example, Duan et al (2001) have shown that aggregation of a model seawater-humic acid solution with FeCl3 occurs at near neutral pH values. This was studied by monitoring floc size, solution pH, and zeta potential. By pH adjustment to 6, the greatest humic acid removal (by coagulation and subsequent membrane filtration) and the largest floc size was achieved at a FeCl3 dosage of 200 µmol/L. It is believed that the coagulation is characterized by competition between OH- ions and humic acid for ferric ions in the co-precipitation process producing hydroxides.

The importance of pH is further stressed in work by Deiana et al (1995). The reduction of ferric iron by natural humic acid was studied in aqueous solution as a function of pH, time and ferric iron concentration. The information gained from spectroscopy (Fourier Transform-IR and electron spin resonance spectroscopy) as well as potentiometric data suggests that redox reactions occur at a low pH due to the involvement of phenolic groups and radicals. At pH values higher than 3.5 the reaction was strongly inhibited by the formation of iron (III)-humate complexes.


In summary, in the environment, a number of important steps follow from any releases. In effect, ferrous and ferric ions can be treated together, because the ferrous ion is rapidly transformed to ferric ion under the conditions found at typical points of release, where the half-times are considered clearly below 12 h.

The iron II & III ion stability is depending on hydrolytical, oxidative and light-induced reductive speciation processes determining the physical form to which the ions are transformed. The rate of transformation is determined by pH, ionic strength of the aqueous medium, the anions present in solution such as sulphate and chloride, the temperature, the partial pressure of oxygen, the initial concentration of ferrous kation, the presence of NOM and humic substances and last but not least light intensity.

Thus it can be concluded that under (common) environmental conditions relevant for risk assessment iron is considered stable as a little dissolute kations in water in hydrolysis and redox equilibrium with ferric hydroxide and oxide complexes, which are very insoluble and largely confined to particulate phases. In soils and sediments these precipitated materials accumulate, polymerize and finally integrate in the matrix.

  • Deiana S, Gessa C, Manunza B, Rausa R, Solinas V (1995). Iron(III) reduction by natural humic acids: A potentiometric and spectroscopic study, European Journal of Soil Science 46(1):103-8.
  • Duan J, Graham NJD, Wilson F (2003). Coagulation of humic acid by ferric chloride in saline (marine) water conditions, Water Science and Technology 47(1, Asia Environmental Technology 2001):41-4.
  • ECHA European Chemicals Agency (2011). Guidance on the Application of the CLP Criteria. Self-published by ECHA Reference: ECHA-11-G-06-EN, Date: 04/2011. 491 p.
  • Hilli T, Pienimäki M (2003). Oulun edustan vesistötarkkailu v. 2003. Jaakko Pöyry Infra. 9M030197. 41 p.+ annexes.
  • IHCP, DG Joint Research Centre, European Commission (2009). Guidance to Regulation (EC) No 1272/2008 on Classification, Labelling and Packaging of substances and mixtures. 550 p.
  • Johnson I, Sorokin N, Atkinson C, Rule K, Hope S-J (2007). Proposed EQS for Water Framework Directive Annex VIII substances: iron (total dissolved). ISBN: 978-1-84432-660-0. Science Report: SC040038/SR9. SNIFFER Report: WFD52(ix). Product Code SCHO0407BLWB-E-E. Self-published by Environment Agency, Almondsbury, Bristol BS32 4UD, U.K. 65 p.
  • Kappler A, Benz M, Schink B, Brune A (2004). Electron shuttling via humic acids in microbial iron(III) reduction in a freshwater sediment, FEMS Microbiology Ecology 47(1):85-92.
  • Morel FMM, Hering JG (1993). Principles and Applications of Aquatic Chemistry, published by Wiley-IEEE.
  • Schnitzer M (1969). Reactions between fulvic acid, a soil humic compound, and inorganic soil constituents. Soil Science Society of America Proceedings 33:75–81.
  • Schwarzenbach RP, Gschwend PM, Imboden DM (1993). Environmental organic chemistry. Wiley Interscience, ISBN 0-471-83941-8, p 266 et seq.
  • Skeaff JM (2004). Review of the Oxidation of Ferrous Ion in Aqueous Media. Work Performed for Arcelor (Luxembourg, Luxemburg), CEFIC (Brussels, Belgium), EUROFER (Brussels, Belgium), Rio Tinto plc. (London, SW1Y 4LD, U.K.). Rev 10; Aug 17 2004. Self-published CANMET-MMSL 04-035 (CR)/Contract No. 602866 Natural Resources Canada. August 2004. 34 p.
  • Tratnyek PG, Macalady DL (2000). Oxidation-reduction reactions in the aquatic environment. IN: Handbook of Property Estimation Methods for Chemicals, 383-415. Boethling RS, Mackay D (Eds). Lewis Publishers, Boca Raton, FL, U.S.A.
  • Wetzel RG (1983). Limnology. 2nd Edition; Complete Revision. Saunders College Publishing, Philadelphia. 858 pp.