Ferrihydrite 349 solubility

This simplified calculation is used to illustrate basic computational techniques. It assumes that all of the Fe(OH)3(aq) is a true solute. The quality of this assumption is a matter of debate as at pH 8, Fe(OH)3(aq), tends to form colloids. Thus, laboratory measurements of ferrihydrite solubility yield results highly dependent on the method by which [Fe(lll)]jQ(gj is isolated. Ultrafiltration techniques that exclude colloids from the [Fe(lll)]jQjgj pool produce very low equilibrium solubility concentrations, on the order of 0.01 nM. This is an important issue because a significant fraction of the iron in seawater is likely colloidal, some of which is inorganic and some organic. In oxic... 

Typically, mammalian ferritins can store up to 4500 atoms of iron in a water-soluble, nontoxic, bioavailable form as a hydrated ferric oxide mineral core with variable amounts of phosphate. The iron cores of mammalian ferritins are ferrihydrite-like (5Fe203 -9H20) with varying degrees of crystallinity, whereas those from bacterioferritins are amorphous due to their high phosphate content. The Fe/phosphate ratio in bacterioferritins can range from 1 1 to 1 2, while the corresponding ratio in mammalian ferritins is approximately 1 0.1. 

As we pointed out earlier, the H subunit catalyses the ferroxidase reaction, which occurs at all levels of iron loading, but decreases with increasing amounts of iron added (48-800 Fe/ protein). Reaction (19.8) catalysed by both FI- and L-chain ferritins, occurs largely at intermediate iron loadings of 100-500 Fe/protein. Once nucleation has taken place, the role of the protein is to maintain the growing ferrihydrite core within the confines of the protein shell, thus maintaining the insoluble ferric oxyhydroxide in a water-soluble form. 

Figure 12. Possible isotope fractionation steps during anaerobic photosynthetic Fe(II) oxidation (APIO). It is assumed that the process of oxidation proceeds through an oxidation step, where Fe(II),q is converted to soluble Fe(III) in close proximity to the cell, followed by precipitation as ferric oxides/hydroxides. As in DIR (Fig. 5), the most likely step in which the measured Fe isotope fractionations are envisioned to occur is during oxidation, where isotopic exchange is postulated to occur between pools of Fe(II) and Fe(III) (Aj). As discussed in the text and in Croal et al. (2004), however, it is also possible that significant Fe isotope fractionation occurs between Fe(III), and the ferrihydrite precipitate (Aj) in this case the overall isotopic fractionation measured between Fe(II), and the ferrihydrite precipitate would reflect the sum of A and Aj, assuming the proportion of Fe(III) is small (see text for discussion). Isotopic exchange may also occur between Fe(II),q and the ferric hydroxide precipitate (Aj), although this is considered unlikely.

If the process of APIO is properly described by Equation (19), which infers the presence of a soluble Fe(III) intermediate species, it will be difficult to analyze this species directly, given the low levels that are expected. We must therefore develop mathematical approaches to estimating the isotopic composition of this component, as was done for DIR. The equations used in the previous chapter (Chapter lOA Beard and Johnson 2004) to describe abiotic Fe(II) oxidation are useful for illustrating possible isotopic fractionations that may occur during APIO. We will assume that the overall oxidation process occurs through a series of first-order rate equations, where relatively slow oxidation of FefTI) to a soluble Fe(III) component occurs, which we will denote as Fe(III)jq for simplicity. The oxidation step is followed by precipitation of Fe(III)jq to ferrihydrite at a much faster rate, which maintains a relatively low level of Fe(III)jq relative to Fe(II)jq. The assumption of first-order kinetics is not strictly valid for the experiments reported in Croal et al. (2004), where decreasing FefTI) contents with time do not closely follow either zeroth-, first-, or second-order rate laws. However, use of a first-order rate law allows us to directly compare calculations here with those that are appropriate for abiologic Fe(II) oxidation, where experimental data are well fit to a first-order rate law (Chapter lOA Beard and Johnson 2004). 

Since [Fe(lll)]jojaj [Fe " ], the formation of ion pairs and complexes is greatly enhancing the equilibrium solubility of ferrihydrite. This is called the salting-in effect and illustrates why mineral solubility calculations in seawater must take ion speciation into consideration. 

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. 

Solubility diagrams have nearly always been calculated using solubility and stability constants. Experimental determination of the solubility of iron oxides as a function of pH has been concerned predominately with ferrihydrite. Lengweiler et al. 

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 equilibrium solubility of an Fe oxide can be approached from two directions -precipitation and dissolution. The first method involves precipitating the oxide from a supersaturated solution of ions with stepwise or continuous addition of base und using potentiometric measurements to monitor pH and calculate Fej- in equilibrium with the solid phase until no further systematic change is detected. Alternatively the oxide is allowed to dissolve in an undersaturated solution, with simultaneous measurement of pH and Fejuntil equilibrium is reached. It is essential that neither a phase transformation nor recrystallization (formation of larger crystals) occurs during the experiment this may happen with ferrihydrite which transforms (at room temperature) to a more condensed, less soluble phase. A discussion of the details of these methods is given by Feitknecht and Schindler (1963) and by Schindler (1963). 

In soils and other geoenvironments, water soluble humic compounds are also candidates for complexation and release of Fe from Fe oxides and thereby may provide Fe for plants. For example, compared to the control, water extractable humics from peat (1.7 mmol C/L) doubled the amount of Fe extracted from a freshly prepared 2-line ferrihydrite over 24 hr (Cesco et al. 2000). 

They are stable only at low redox potential. They form either by direct precipitation from an Fe" salt solution upon oxidation once their solubility product is exceeded (eqn. 13.13), or by interaction between 2-line ferrihydrite precipitated initially and Fe " in solution (eqn. 13.14) in the presence of a sufficiently high [Fe " ], the green rust is more stable than 2-line ferrihydrite... 

The transformation has been followed up by XRD, Mbssbauer spectroscopy, EXAFS and colorimetry. It can be monitored more conveniently, however, by the acid oxalate extraction method in which residual ferrihydrite is dissolved and the crystalline product left intact (Schwertmann Fischer, 1966). The extent of transformation at any time is given as the ratio FOo/Fet where Fe is the oxalate soluble iron (i. e. the unconverted ferrihydrite) and Fet is the total iron in the system. A plot of log (FCo/Fet) against time of aging at 100 °C is linear over 90-95 % of the reaction... 

A 6-line ferrihydrite whose water content was reduced from 146 to 26 g kg by heating for 3000 hr at 123 °C, while the oxalate soluble proportion decreased from 100% to 12% and the unit cell volume from 0.3091 to 0.3079 nm ... 

The foreign species act in solution and usually retard nucleation or growth of goethite by competing with soluble Fe " species for sites on the subcritical nucleus or on the growing crystal. This mechanism is independent of the presence of ferrihydrite. 

Fig.14.20 Effect of various clay minerals on the transformation of 2-line ferrihydrite to goethite and hematite at 25 °C and pH 5 after 16 yr as measured by the ratio of oxalate to dithionite soluble Fe (Feo/Fed) (Schwertmann et al.

With M/(Fe + M) >0.15, a spinel phase (MFe204) formed in all cases and when this ratio exceeded 0.33, Cu and Ni precipitated as separate phases (Tab. 14.3). Formation of a spinel phase requires a threshold level of in the system. It is considered that the spinel phase nucleates in the water layer adsorbed on or adjacent to, the surfaces of the ferrihydrite particles and that these nuclei grow by addition of soluble M-Fe-hydroxo complexes released by the dissolving M-ferrihydrite (Cornell Giovanoli, 1987, 1989 Giovanoli Cornell, 1992). Tronc et al (1992) suggested that when the ratio is very low, a different mechanism operates a mixed valence... 

Schwertmann, 1993). Such soils are characterized by a hydraulic conductivity somewhere in the profile which is too low to cope with the high rainfall, so that all pores will be filled with water for certain periods of time (see above). In this case, the oxygen supply is limited by the low level of O2 dissolved in the soil water (46 mg O2 at 25 °C) and reduction of Mn-oxides, nitrate and Fe oxides sets in. Soils containing Fe oxides are, therefore, redox-buffered (poised). The redox titration curve (Fig. 16.14) of a soil with 23 g kg Fe as Fe oxides shows buffering at two different pe -1- pH levels, one at ca. 11 and another at ca. 9, which indicate the presence of a more reducible (e. g. ferrihydrite) and a less reducible (e. g. goethite) Fe oxide, respectively, in accordance with their different solubilities (see Chap. 9). 

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