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EC number: 231-843-4
CAS number: 7758-94-3
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
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
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).
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
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).
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
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.
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).
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
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
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
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
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