Reaction mass of aluminium chloride, iron dichloride, magnesium chloride and manganese dichloride

Endpoint summary

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.

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.

Hydrolysis

Ferric ions released into (or generated in) water will rapidly precipitate as highly insoluble oxides and oxo-hydroxides. These stable compounds are exactly the forms in which iron is found naturally in the earth’s crust.

Phototransformation

Iron in the ferrous form is subject to phototransformations. This mechanistically-complex process is well-known in the general scientific literature, and has been reported in respect of the speciation of marine iron. Ferrous sulphate and ferrous chloride are reducing agents, and exposure to light gives reduction of the ferric form.

Reactions with oxygen

In water containing oxygen, the chemistry is dominated by iron-oxygen reactions. These are highly complex mechanistically (Cotton and 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). 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 (Kopacek 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 litre-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:

• is highly dependent on pH (100 fold 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 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 hr, 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. 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 hr and 90% removal would be achieved by 36 hr. 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).

• 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.

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

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.

• 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.
• Hilli T, Pienimäki M (2003). Oulun edustan vesistötarkkailu v. 2003. Jaakko Pöyry Infra. 9M030197. 41 p.+ annexes.
• 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 & Hering 1993
• Morel FMM, Hering JG (1993). Principles and Applications of Aquatic Chemistry, published by Wiley-IEEE
• Schwarzenbach RP, Gschwend PM, Imboden DM (1993). Environmental organic chemistry. Wiley Interscience, ISBN 0-471-83941-8, p 266 et seq.
• Wetzel, R. G. 1983. Limnology. (2nd Edition; Complete Revision) Saunders College Publishing, Philadelphia. 858 pp

Manganese

Redox conditions and the pH level have a determinative influence on the manganese speciation in soils, sediments and waters. The following overview is in accordance to the respective sections of the CICAD 63 (WHO 2004 and 2005). There is little evidence for manganese–organic associations in natural waters, with manganese only weakly bound to dissolved organic carbon (L’Her Roux et al 1998). Hence, organic complexation does not play a major role in controlling manganese speciation in natural waters. Organomanganese compounds like Methylcyclopentadienyl Manganese Tricarbonyl (MMT, CAS 12108-13-3) are known. They were industrially produced but their formation under environmental conditions is unlikely.

Aquatic

While at lower pH and redox potential values manganese (II) is predominantly present, the percentage of oxyhydroxide species in colloidal form becomes significant above pH 5.5 in moderately eutrophic waters (LaZerte & Burling 1990). In aerobic environments manganese (II) is subject to oxidation and consequent precipitation and adsorption. Eventually manganese dioxide (CAS 1313-13-9) is formed, which is insoluble and thus biologically unavailable. But at pH < 8.5 the kinetics of these processes are slow (Zaw & Chiswell 1999). The time required is in the order of days in natural waters (Stokes et al 1988). Generally the oxidation rates increase with increasing pH or the presence of catalytic surfaces such as manganese dioxide (Huntsman & Sunda, 1980), which is itself an oxidation product and indicates an autocatalytic tendency. In the sequence of the complex oxidation reactions to the final precipitation as manganese dioxide several species, i.e. dissolved Mn(II), hydrous oxides of Mn(III), Mn(II) adsorbed to particulates, and Mn(II)–ligand complexes occur. Additional factors influencing these speciations are inorganic and organic carbon, sulphate, chloride, temperature, and time (Stokes et al 1988). In consequence the speciation in any particular river or stream may be different depending principally on the hydrogeological conditions of the catchment at time of sampling. The suspended sediment will be mixed with soluble manganese (II) species in varying proportions.

Terrestrial

Water-soluble manganese in soils is directly proportional to the pH value as oxidation kinetics depends on the latter. Due to the dependence of the solubility from the oxidation state soluble manganese (II) predominates at lower pH. Reducing conditions will result in higher concentrations of bioavailable dissolved manganese, e.g. in flooded soils (Stokes et al 1988). However, there is competition by iron, and plant absorption of manganese is decreased or unaffected by flooding (Adriano 1986).

According to Evans (1989) there are two main mechanisms involved in the retention of manganese by soil.

1. The charged surface of soil particles forms with the manganese ions manganese oxides, hydroxides, and oxyhydroxides through kation exchange reactions. These manganese species represent adsorption sites for other metals.
2. Adsorption of manganese to other oxides, hydroxides, and oxyhydroxides result from ligand exchange reactions.

• Adriano DC (1986). Trace elements in the terrestrial environment. New York, NY, Springer-Verlag.
• Evans LJ (1989). Chemistry of metal retention by soils: Several processes are explained. Environmental Science & Technology, 23:1048–56.
• Huntsman SA, Sunda WG (1980). The role of trace metals in regulating phytoplankton. In: Morris I, ed. The physiological ecology of phytoplankton. Berkeley, CA, University of California, pp. 285–328.
• L’Her Roux L, Le Roux S, Appriou P (1998) Behaviour and speciation of metallic species Cu, Cd, Mn and Fe during estuarine mixing. Marine Pollution Bulletin, 36(1):56–64
• LaZerte BD, Burling K (1990). Manganese speciation in dilute waters of the Precambrian Shield, Canada. Water Research 24:1097–101.
• Stokes PM, Campbell PGC, Schroeder WH, Trick C, France RL, Puckett KJ, LaZerte B, Speyer M, Hanna JE, Donaldson J (1988). Manganese in the Canadian environment. Ottawa, Ontario, National Research Council of Canada, Associate Committee on Scientific Criteria for Environmental Quality (NRCC No. 26193).
• WHO World Health Organization (2004 and 2005). Manganese and its Compounds: Environmental Aspects. Concise International Chemical Assessment Document 63, Corrigenda published by 12 April 2005 have been incorporated. Self-published, Geneva, Switzerland
• Zaw M, Chiswell B (1999). Iron and manganese dynamics in lake water. Water Research 33(8):1900–10.

Aluminium

Aluminium cannot be degraded in the environment as it is an element, but may undergo various precipitation or ligand exchange reactions (ATSDR 2008). Aluminium is the most abundant metal in the lithosphere, and is characterized by a complex biogeochemical cycle (Driscoll & Postek 1996, Exley 2003). Nonetheless aluminium has in compounds only one oxidation state (+3), and would under environmental conditions not undergo oxidation-reduction reactions (ATSDR 2008). Aluminium can participate in hydrolysis reactions, thereby forming a number of monomeric and polymeric Al-hydroxides and this process is highly dependent on pH. Aluminium can be complexed by various ligands present in the environment, e.g. fulvic and humic acids. The solubility of aluminium in the environment will depend on the ligands present and the pH (ATSDR 2008).

• ATSDR Agency for Toxic Substances and Disease Registry (2008). Toxicological Profile for Aluminum. U.S. Department of Health and Human Services, Public Health Service, Atlanta, Georgia, U.S.A. 357 p
• Driscoll CT, Postek KM (1996). The chemistry of aluminium in surface waters. IN: Sposito G (Ed) The environmental chemistry of aluminium. 2nd edition. Boca Raton (FL, U.S.A.) CRC Press. p 363-418.
• Exley C (2003). A biogeochemical cycle for aluminium. Journal of Inorganic Biochemistry 97:17.