Registration Dossier

Environmental fate & pathways

Hydrolysis

Currently viewing:

Administrative data

Link to relevant study record(s)

Description of key information

The behaviour of Fe2+ and Fe3+ metal ions in solution is dependent on the different conditions. Key conditions that influence iron behaviours are the oxygen content, pH, and the presence of potential ligand anions with which the kationic iron might associate. Hydrolysis is relevant for iron and forms precipitating hydroxylates. A hydrolysis value of 70 % and 90 % removal from solution is calculated to be achieved in 19 hour and 36 hour respectively

Key value for chemical safety assessment

Additional information

Irrelevance of hydrolysis as decomposition process

Since hydrolysis changes the chemical form but does not decompose metal species and since characterization of total metal concentration considers all chemical forms, the concept of degradation of metals by hydrolysis is not relevant in the consideration of their environmental fate. Nonetheless several hydrolysis reactions of metal kations are known and of importance as the formed hydroxides play an important role with regard to precipitation and bioavailability.

Oxidation of iron (II) under environmental conditions forms iron (III) rapidly

Iron (II) ions may be oxidised to iron (III) under most environmental conditions. The iron (II) ion can be oxidised by common oxidants such as nitrate (Johnson et al 2007 with reference to Schnitzer 1969). Ferrous ion (Fe II) is unstable when exposed to air, it oxidised to ferric ion (Fe III), which in turn forms the insoluble, hydrated, amorphous and gelatinous precipitate of ferric hydroxide Fe(OH)3. The expected rapid oxidation of iron (II) to iron (III) is the precondition to treat their hydrolysis behaviour together and to consider them a category.

Hydrolysis of Iron (III)

Johnson et al (2007) summarize the hydrolysis of iron (III) with reference to Schnitzer (1969) as follows: “In solution, iron(III) ions are expected to hydrolyse or form complexes. At pH <1, the hexa-aqua ion ([Fe(H2O)6]3+) is the predominant species. As the pH increases above 1, a stepwise hydrolysis occurs. Between pH 1–2, various species of hydroxo- and oxo-iron compounds may be formed. Above pH 2, colloidal gels are formed, giving a precipitate of the red–brown gelatinous hydrous iron oxide. In the presence of complexing anions, such as chloride, the hydrolysis of iron (III) can result in chloro-, aqua- and hydroxo-species.”

Iron ions at equilibrium in water

When ferric ion is added to water the hexa-aquo kation is formed. This is strongly acidic with a pKa of 3.05 (Cotton & Wilkinson 1972). Thus:

[Fe(H2O)6]3+ → [Fe(H2O)6](OH)2+ + H+(aq)

The complete hydrolysis of Fe(III) follows the reaction:

Fe3+ + 3 H2O <=> Fe(OH)3 (s) + 3 H+

The ferrous Fe(II) ion stable only under non-oxygenated conditions and is not acidic in solution. The importance of pH is further emphasised by consideration of the solubility product (Ksp) values of the hydroxides. The equations defining solubility product are:

Ksp = [Fe2+][OH-]2 for ferrous hydroxide, and

Ksp = [Fe3+][OH-]3 for ferric hydroxide.

Thus ferrous ion Fe(OH)2, can be formed, it is moderately insoluble, with

Ksp = 1.6 x 10^-14 (Lide 2009).

Ferric hydroxide (Fe(OH)3) is highly insoluble with Ksp = 1 x 10^-36 (Lide 2009). Formation of ferric hydroxide at pH levels above 5.0 limits the presence of iron in aqueous systems.

The significance of pH on the solubility of ferrous and ferric can be seen in the respective Table in the section on water solubility. The implication of this analysis is that under conditions of very low oxygen concentration, ferrous is freely soluble but ferric is not. Under conditions of high concentration and low oxygen, ferric ion could acidify the water, thereby having environmental consequences. These conditions would not apply in the normal direct uses, although could occur during major accidental leakage.

The potential of the Fe(III)-Fe(II) couple, (0.77 V) is such that molecular oxygen can convert ferrous to ferric in acid solution or basic solution (Cotton & Wilkinson 1972).

An in-depth analysis of the oxidation and precipitation of iron was carried out by CEFIC as part of the recent European Chemicals Bureau classification process of ferrous sulphate (Skeaff 2004).

In conclusion at environmental pH (5-9, ECHA 2012, p 69 ), 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.

  • Cotton FA, Wilkinson G eds (1992). Advanced Inorganic Chemistry. John Wiley and Sons, New York, U.S.A.
  • 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.
  • ECHA European Chemicals Agency (2012). Guidance on information requirements and chemical safety assessment Chapter R.7a: Endpoint specific guidance. First Draft for PEG 2nd March 2012. Version 2. Guidance for the implementation of REACH. Self-published March 2012. 167 p.
  • Lide DR ed (2009). CRC Handbook of Chemistry and Physics. 90th print run. Taylor & Francis, ISBN 978-1-4200-9084 -0
  • Schnitzer M (1969). Reactions between fulvic acid, a soil humic compound, and inorganic soil constituents. Soil Science Society of America Proceedings 33:75–81.