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

Hydrolysis

Description of key information

The behaviour of metal ions, e.g. iron (Fe2+ and Fe3+) 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 cationic iron might associate. Hydrolysis is relevant for iron and aluminium and forms precipitating hydroxylates.

Additional information

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.

Iron

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.

This discussion is set out stepwise in respect of the equilibrium state, reaction rates, and conclusions. It is applied further in the health and environmental effects sections.

Iron ions at equilibrium in water

Under non-oxygenated conditions, when ferric ion is added to water the hexa-aquo kation is formed. This is strongly acidic with a pKa of 3.05 (Cotton and 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 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 Table below. This shows the maximum dissolved concentration of the metal ion in pure water as it depends upon pH.

Table: Calculated maximum solubility of iron in solution at pH range of 4.0 to 8.0 and temperature of 20 °C:

	Ferrous		Ferric
pH	Fe (mg/L)	mmol/L	Fe (mg/L)	mmol/L
4	>1E+06	>1.8E+04	6.16E-02	1.1E-03
5	>1E+06	>1.8E+04	6.16E-05	1.1E-06
6	>1E+06	>1.8E+04	6.16E-08	1.1E-09
7	89600	1600	6.16E-11	1.1E-12
8	896	16	6.16E-14	1.1E-15

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 and 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). A review of the scientific literature on the oxidation of ferrous sulphate reveals the following:

Ferrous sulphate reacts with water to form ferrous hydroxide (Fe(OH)2), moderately insoluble, Ksp = 1.6 x 10^-14) (Lide 2009, Cotton and Wilkinson 1972). Any precipitate would in turn undergo further oxidation to form ferric hydroxide (Fe(OH)3) which is highly insoluble (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.

• Lide DR (Ed) 2009 CRC Handbook of Chemistry and Physics. 90th print run. Taylor & Francis, ISBN 978-1-4200-9084-0

Manganese

However manganese undergoes speciation in the environment the formation of hydroxides and subsequent precipitation are no relevant processes in the environment.

Aluminium

Aluminium is the most abundant metal in the lithosphere, and is characterized by a complex biogeochemical cycle (Driscoll and Postek 1996, Exley 2003). Aluminium can participate in hydrolysis reactions, thereby forming a number of monomeric and polymeric Al-hydroxides and this process is highly dependent on pH. The aluminium ion is surrounded by six water molecules in solution (Cotton et al 1999). The hydrated trivalent ion, i.e. [Al(H2O)6]3+, undergoes hydrolysis. The coordinated water ligands get stepwise deprotonated and thus transformed to hydroxide ligands. Thus e.g. the following species are produced: [Al(H2O)5(OH)]2+, [Al(H2O)4(OH)2]+ (Snoeyink & Jenkins 1980). Aluminium persists in the environment irrespective of whatever chemical species form as a result of hydrolysis, although it may form insoluble aluminium hydroxides that precipitate out of solution.

The aluminium speciation in water depends on pH (ATSDR 2008). The predominant form at pH levels below 4 is[Al(H2O)6]3+. Hydrolysis at pH 5 to 6, reveals Al(OH)2+ and Al(OH)2 +. In the pH range from 5.2 to 8.8 Al(OH)3 is formed predominantly. Above pH 9 the species Al(OH)4, which is soluble dominates and is the one present species above pH 10 (Martell & Motekaitis 1989).

Between pH 4.7 to 10.5 polymeric aluminium hydroxides occur. They tend to increase in size until from eventually amorphous colloidal particles, which consist in sum of Al(OH)3. The latter one crystallise to gibbsite in acid waters (Brusewitz 1984). Polymerization is affected by the presence of dissolved silica; when enough silica is present, aluminium is precipitated as poorly crystallized clay mineral species (Bodek et al 1988).

Characterization of aluminium in environmental media is typically based on total aluminium concentrations inclusive of all specific chemical forms or species.

Hydroxyaluminium compounds are considered amphoteric, e.g. they can act as both acids and bases in solution (Cotton et al 1999). Because of this property, aluminium hydroxides can act as buffers and resist pH changes within the narrow pH range of 4–5 (Brusewitz 1984).

• Brusewitz S (1984). Aluminum. Vol. 203. Stockholm, Sweden. University of Stockholm, Institute of Theoretical Physics.
• Bodek I, Lyman WJ, Reehl WF, et al (Eds) (1988) Aluminum (Al). Environmental inorganic chemistry: Properties, processes, and estimation methods. New York, NY, U.S.A., Pergamon Press, 6.7-1 to 6.7-9.
• Cotton FA, Wilkinson G, Murillo CA, et al (Eds) (1999). The group 13 elements: Al, Ga, In, Tl. Advanced inorganic chemistry. 6th ed. New York, NY, U.S.A., John Wiley & Sons, p 175-207.
• 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.
• Martell AE, Motekaitis RJ (1989). Coordination chemistry and speciation of Al (III) in aqueous solution. IN: Lewis TE (Ed) Environmental chemistry and toxicology of aluminum. Chelsea, MI, U.S.A., Lewis Publishers, p 3-17.
• Snoeyink VL, Jenkins D (Eds) (1980). Water chemistry. New York: John Wiley and Sons, 146, 209-210.