Goethite solubility

In the ultimate analysis it may be pointed that the aforesaid hydrolysis processes are no doubt technically very satisfactory and tolerable, but environmentally this is not the case. The different processes yield jarosite, goethite and hematite, all of which retain considerable amounts of other elements, especially, zinc and sulfur. The zinc originates mainly from undissolved zinc roast in the iron residues, and sulfur from sulfate, which is either embodied into the crystal lattice or adsorbed in the precipitate. As a consequence of the association of the impurities, none of these materials is suitable for iron making and therefore they must be disposed of by dumping. The extent of soluble impurities present in the iron residues means that environmentally safe disposal not an easy task, and increasing concern is being voiced about these problems. An alternative way of removing iron from... 

The primary anion studied in both the titration calorimetry7 and CIR-FTIR experiments reported here was the salicylate (2-hydroxybenzoate) ion (SAL). Acidity constants for salicylic acid are pK = 3.0 and pK = 13 (12), and the aqueous solubility of salicylic acid is 2.4 g/L, while that of NaSAL is 975 g/L. SAL has been shown to adsorb on both iron oxides (13) and aluminum oxides (14). Several other anions were also studied, and results for these anions are given as needed to illuminate certain features of the salicylate-goethite adsorption process. 

In the sorption experiments of Icopini et al. (2004), the measured isotopic contrast between Fe(II)aq and the goethite starting material was -0.8%o after Fe(II) had sorbed to the surface over 24 hours in this case, the isotopic fractionation between sorbed Fe(II) and Fe(II)aq is not the 0.8%o measured difference, but is approximately +2.1%o based on an inferred 8 Fe value for the sorbed component as calculated from Fe mass balance (Fig. 4), as was noted in that study. Measured differences in Fe isotope compositions between ferric oxide/hydroxide and Fe(II)aq during dissimilatory Fe(III) reduction and photosynthetic Fe(II) oxidation have been proposed to reflect fractionation between soluble Fe(II) and Fe(III) species, where the soluble Fe(III) component is postulated to be bound to the cell and is not directly measured (Beard et al. 2003a Croal et al. 2004). In the case of dissimilatory Fe(III) reduction, assuming a static model simply for purposes of illustration, if 50% of the Fe in a pool that is open to... 

The extent to which a sparingly soluble solid dissolves is expressed by the solubility product. This describes the equilibrium established between the solid and the concentration of its ions in a saturated solution. Consider, for example, the dissolution of goethite in water ... 

The sum of all the soluble Fe " species, i. e. Fer, in equilibrium with goethite as a function of pH, is the heavy line in the solubility diagram in Figure 9.1, whereas the activities of the single species are shown by the weak lines. Inclusion of the hydrolysis species results in a much higher solubility than would be observed by consideration of the solubility product i. e. Fe ", alone. For example, at pH 6, is <10 M, whereas Upe. = 10 M. Only at very low and very high pH is Up. essentially equal to UpeJt and ape(OH)j. respectively. 

The solubility plots for lepidocrocite, ferrihydrite and hematite (Fig. 9.2) and for goethite, ferrihydrite and soil-Fe (Fig. 9.3) show only the total Fe activity. They were obtained in the same way as that for goethite using the appropriate constants from Tables 9.1, 9.2 and 9.4. 

Fig. 9.3 Solubilities of goethite, ferrihydrite and soil-Fe as a function of pH (data for Soil Fe from Lindsay, 1979, with permission).

Fig. 9.4 Calculated and experimental solubility of ferrihydrite as a function of pH (Schindler et al., 1963, with permission). The curves were calculated taking into account the species Fe FeOH Fe(OH)J, Fe(OH)4 and Fe2(OH) and the following solubility products (log - Kso) ferrihydrite, freshly precipitated 3.96 ferrihydrite, aged 3.55 and goethite 1.4.

Lengweiler et al. (1961) found that the solubility of goethite, like that of ferrihydrite, increased as the pH rose above 12. For ferrihydrite, equilibrium between the solid and Fe(OH)4 was reached quite rapidly, whereas for goethite, equilibrium was not reached even after 40 days (25 °C). A value of 1.40 + 0.1 for of goethite (surface area ca. 100 m g" ) was only reached after 3 years (Fig. 9.5) (Bigham et al., 1996). As expected on thermodynamic grounds, the solubility of goethite was 10 to 10 times less than that of ferrihydrite. 

The effect of reducing conditions on the solubility of an iron oxide can be found by combining the appropriate dissolution equations with the redox equation to obtain the concentration of the Fe" species released. In the Fe"/Fe oxide system, protons are always involved because the state of hydrolysis of the Fe is changed. For goethite, for example. 

Fig. 9.7 The solubility of goethite under oxidizing and reducing conditions at (pe + pH) = 8.

Fig. 9.8 Particle size effect on the solubility products of goethite and hematite (Langmuir and Whittemore, 1971, with permission).

In this case, d (nm) is the diameter of spherical particles. Both sets of equations are only approximations, particularly for goethite, the particles of which are often acicular. However, they do enable an estimate of the rise in solubility, as particle size drops, to be obtained (see Fig. 8.3). There is little difference between the results calculated using the two sets of equations for particles >100 nm, but for 10 nm particles there is more than an order of magnitude difference between the two equations. The higher solubility of smaller particles may lead to their transformation to larger ones via solution, a process called Ostwald ripening. 

Solubility products may be calculated using the free energies of formation of the oxide, the free metal ion, OH and water (Tab. 9.1 and 9.2). For goethite, for example. 

Researchers in the aluminium industry have investigated the solubility of goethite in sodium aluminate and NaOH solutions. Basu (1983) found, using samples of natural goethite, that the equilibrium solubility of goethite in sodium aluminate solution was close to zero at room temperature and increased exponentially as the temperature rose above 100 °C. She also found that the isothermal solubility was greater in 5 M NaOH than in 5 M sodium aluminate solution at 150 °C, for example, [Fej] was 20 and 50 mgL , respectively. 

Interaction of phosphate solutions with goethite may lead to surface precipitation of phosphates if the concentration of P in solution exceeds the mineral solubility (Jo-nasson et al., 1988). A combined Auger, XPS, scanning SIMS and electron diffraction study showed that after 90 days at 60 °C, crystals of griphite (an Fe hydroxy phosphate) precipitated out of a phosphate solution onto crystals of goethite (Martin et al., 1988). 

Goethite forms in aqueous media by direct precipitation from soluble Fe " species which are supplied by hydrolysis of Fe " solutions, by dissolution of a solid precursor, or by oxidation/hydrolysis of Fe" salt solutions. 

In moderately alkaline solutions (pH >8) oxidation of Fe" solutions proceeds via Fe(OH)2 and usually yields magnetite (David Welch, 1956 Sidhu et al., 1977). Under these conditions the solubility product of magnetite is exceeded so the mixed oxide is more stable than the pure Fe " oxides (see Chap. 8). Tamaura et al. (1981) monitored the transformation of Fe(OH)2 at pH 11 and 65 °C. Initially both goethite... 

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