Mill scale (ferrous metal)

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Description of key information

 

The fate and behaviour of iron in aquatic systems is particularly complex because, under the range of natural conditions found in freshwaters, iron can undergo changes in oxidation state which affects the way in which it behaves. The majority of the iron which enters the aquatic environment from natural sources is mobilised as Fe(II). This is due to the much greater solubility of Fe(II) compared to Fe(III), although this iron is transformed to Fe(III) by oxidation processes in oxic aquatic systems. Some iron species are highly insoluble in circumneutral, oxygenated waters resulting in the formation of mineral precipitates. In addition to this behaviour the different forms of truly dissolved iron are also able to bind to organic matter, which may affect their potential for transport and tendency to undergo redox or precipitation changes. It is also particularly difficult to identify the particular forms of iron present in a natural water sample using routine analytical methods such as performing a total metal analysis on filtered samples. Equilibrium speciation calculations may also need to be treated with caution as an equilibrium situation may not exist under natural conditions due to the dynamics of iron chemistry.

 

The soluble or dissolved iron is usually defined operationally as the iron content of water after filtration at 0.45 μm (sometimes 0.22 μm). The “dissolved iron” samples, however, may contain an appreciable concentration of very fine colloidal iron precipitates which are able to pass through a 0.45 μm filter membrane. Sigg et al. (2000) defined this fraction as aquatic particles with a size < 10 μm that do not coalesce rapidly into bigger particles. These forms of iron would not normally be considered to be in true solution, but are unresolvable from the truly dissolved forms of iron in the filtration and analysis.

In practice, dissolved iron concentrations represent ‘free’ iron, hydrolyzed iron, complexed iron, polymeric iron, and fine colloidal iron.

 

Total iron is the iron content of water plus particles, after acid digestion. The use of mineral acids in varying strengths can lead to different degrees of dissolution of the recalcitrant or tightly bound iron particularly in samples with high suspended sediment content. Typical analytical procedures for determination of total/dissolved iron include inductively coupled plasma atomic emission spectroscopy (ICP-AES), or inductively coupled plasma mass spectrometry (ICP-MS), depending on the concentration range of the samples.

 

All of these properties, together with specific physicochemical processes such as colloid-formation or pH-dependent precipitation make an assessment of the fate and behaviour of iron in aquatic systems particularly complex.

  

Oxidation-reduction reactions with iron

Iron can exist in two oxidation states: Fe(II) is the dominant form of iron under reducing conditions, whereas Fe(III) is the dominant form of iron under oxidising conditions. The rate of oxidation of Fe(II) to Fe(III) in oxic aqueous environments largely depends on pH, with shorter half-lives at higher pH (Skeaff, 2004). This redox-speciation, however, is also influenced by the presence of dissolved organic carbon/matter (DOC/DOM) in solution and the ratio of iron to organic carbon (Weber et al., 2006; Gaffney et al. 2008) and may result in a significant proportion of the iron load in natural waters to remain in a reduced form, even when the waters are well oxygenated. The rate of oxidation of Fe(II) to Fe(III) is also affected by other factors, such as temperature, ionic strength, and the partial pressure of oxygen (Skeaff 2004).

Due to this interaction with DOC/DOM, but also because of formation of Fe(II) containing mineral precipitates or input/output fluxes of Fe(II) and Fe(III,) respectively, a true thermodynamic equilibrium may not actually be achieved. Many natural waters are therefore considered to exist as a “steady-state pseudo-equilibrium” with respect to these types of processes. Natural waters under more acidic conditions are likely to be further from thermodynamic equilibrium than higher pH waters. Calculated speciation of Fe in field waters is therefore only meaningful when all relevant processes are adequately incorporated into the modelling. In practice several different scenarios may need to be considered to understand the true situation.

 

Solubility of iron

Fe(III) is extensively hydrolysed in slightly acidic to neutral freshwaters, which can result in the formation of precipitates due to the low solubility of Fe(OH)3. The solubility of freshly formed amorphous iron oxide precipitates is greater than that of more crystalline, aged, precipitates such as hydrous ferric oxide and ferrihydrite (Tipping et al. 2002).

Fe(II) may also form hydroxide precipitates under the conditions of many natural waters, although oxidation to Fe(III) would be expected, but may be slow at low pH. Fe(II) could be bound to DOC, precipitated iron (oxy)hydroxides, or present as part of the precipitated mineral phases, as well as being present in true solution.

Iron can also be bound to organic matter, in a manner similar to that of many other trace metals, and the complexation can be predicted using chemical speciation models such as WHAM (NERC 2001). It is commonly necessary for the concentrations of iron species in solution to be estimated by assuming that they are at equilibrium with a solid phase, such as ferrihydrite. The two iron ions (Fe(II)and Fe(III)) can bind to dissolved organic matter, but differences in their binding affinities to humic and fulvic acids may be attributed principally to their charge. Electrostatic effects result in a much stronger binding of Fe(III) (when compared to Fe(II)) and therefore all dissolved phase Fe(III)would be expected to be bound to organic matter throughout the range of pH which is relevant to natural waters. Fe(II), on the other hand, would be expected to show steadily increasing DOC binding with increasing pH, up to a maximum at around pH 7.8. The precipitation of iron can also be biologically mediated in some aquatic systems.

Overall, iron solubility is expected to be controlled by the formation of insoluble (oxy)hydroxide precipitates (e.g. ferrihydrite). Truly dissolved concentrations of iron in waters which are at, or close to, thermodynamic equilibrium are expected to be extremely low. The precipitates formed represent a potential important binding phase for other metal ions present in the solution.

 

Iron in the sediment compartment

Sufficiently large enough iron precipitates that are formed in oxic waters will settle out of suspension. When buried in the sediment a remobilisation of iron may occur as Fe(II) will be formed under the reducing conditions of the anoxic sediment layer. A cycle of oxidation, precipitation, reduction and dissolution can therefore exist for iron close to the sediment water interface, and may e.g. lead to considerably elevated dissolved iron concentrations in sediment pore water. These concentration levels usually only accounts for a small proportion of the total loading of iron into the waterbody (Davison et al. 1982).

Iron will also be precipitated as insoluble iron sulphide in reducing sediments where sulphate reducing conditions occur.

 

Adsorption of iron to solid particles

Adsorption of iron in freshwaters is expected to be controlled to a large extent by the pH-dependent solution behaviour of iron; under most environmental conditions the majority of iron present in freshwater samples would be expected to be present as particulate material, although in some situations an appreciable proportion of this may be present as very fine colloidal precipitates which are able to pass through a 0.45 μm filter membrane. As a result of this the measured “dissolved iron” concentration is likely to overestimate the truly dissolved forms of iron. High concentrations of DOC may result in specific binding of truly dissolved iron (i.e. Fe(II)and Fe(III)), or stabilisation of fine colloidal mineral precipitates, reducing the rate at which they are removed from solution/suspension by settling

Comparing Fe(III) to Fe(II), the former has a greater affinity for stronger binding sites (e.g. multi-dentate binding sites). Virtually all dissolved phase Fe(III)would thus be expected to bind to organic matter throughout the range of pH which is relevant to natural waters. Fe(II), on the other hand, is expected to show steadily increasing DOC binding with increasing pH, up to a maximum at around pH 7.8.

The binding of iron by DOC may effectively stabilise the forms of iron which are bound, resulting in a reduced tendency for oxidation/reduction or precipitation. In principle the truly dissolved concentration of iron (i.e. Fe(II)and Fe(III)) may be buffered by solid phase precipitated minerals and also by complexation to colloidal organic matter.

In conclusion, the following processes are likely to reduce the precipitation of iron (oxy)hydroxide mineral phases in the presence of high concentrations of organic ligands:

• reducing the free ion activity of Fe(III)


• stabilisation of Fe(II)against oxidation


• dissolution of precipitated iron,


• stabilization of suspensions of colloidal iron precipitates by the presence of organic matter

 

 

 

Additional information

Iron is a commonly occurring metallic element, comprising 4.6% of igneous rocks and 4.4% of sedimentary rocks (Morel and Hering, 1993). The typical iron concentrations in soils range from 20,000 to 550,000 mg/kg (Bodek et al., 1988), 0.01 mg/l in seawater, up to 0.1 – 10 mg/l in fresh water and between 10,000 – 90,000 mg/kg in sediments (Drever, 1982; Hem, 1970; Horne, 1978; Kabata-Pendias and Pendias, 1984; Khalid et al., 1977; Lindsay, 1979; Stumm and Morgan, 1981). Concentrations of iron can vary significantly, even within localized areas, owing to differences in soil types and the presence of other sources. Iron can occur in either the divalent (ferrous or Fe+2) or trivalent (ferric or Fe+3) states under typical environmental conditions. The valence state is determined by the pH and Eh (redox potential) of the system, and the state of combination of iron is dependent upon the availability of other chemicals (e.g., sulphur is required to produce FeS2 or pyrite) (US EPA 2003a).