Solubility of iron

In general, it is fair to state that one of the major difficulties in interpreting, and consequently in establishing definitive tests of, corrosion phenomena in fused metal or salt environments is the large influence of very small, and therefore not easily controlled, variations in solubility, impurity concentration, temperature gradient, etc. . For example, the solubility of iron in liquid mercury is of the order of 5 x 10 at 649°C, and static tests show iron and steel to be practically unaltered by exposure to mercury. Nevertheless, in mercury boiler service, severe operating difficulties were encountered owing to the mass transfer of iron from the hot to the cold portions of the unit. Another minute variation was found substantially to alleviate the problem the presence of 10 ppm of titanium in the mercury reduced the rate of attack to an inappreciable value at 650°C as little as 1 ppm of titanium was similarly effective at 454°C . 

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. 

Eordham, A.W. (1970) Sorption and precipitation of iron on kaolinite. III. The solubility of iron(III) hydroxides predpitated in the presence of kaolinite. Aust. J. Soil Res. 8 107— 122... 

Lui, X. Millero, F.J. (1999) The solubility of iron hydroxide in sodium chloride solutions. Geochim. Cosmochim. Acta 63 3487-3497 Lumsdon, D.G. Evans, L.G. (1994) Surface complexation model parameters for goethite... 

With reference to the discussion of the bimetallic corrosion of iron given in Section 16.1, confirm (a) that the solubility of iron(II) hydroxide is 5.8 /zmol L-1 if OH- is produced along with iron(II) according to reaction 16.5 and the reverse of reaction 16.6 and (b) that the timescale of the oxidation of iron(II) in solution at pH 7 is as given following Eq. 16.8. Hint Incorporate [O2] and pH into a rate constant ki for a first-order process, then calculate the half-period h/2 = (ln2)/fci. 

Concrete, reliable values on the solubility of Iron Blue are not recorded in the scientific literature. Based on comparative calculations between the known solubility of Fe(OH)3 on the one hand, and the limit value of the pH stability of Iron Blue on the other hand (pH 10), the approximate solubility of Iron Blue in water can, however, be calculated (see chapter 6.6.2.2.). It amounts to approximately 10"24 g Iron Blue per liter of water, this means that 0.000000000000000000000001 g Iron Blue dissolve in 1,000 g of water. 

Therefore, they did not determine the solubility of Iron Blue, but the measure of stability of the dispersion of fresh precipitations of the pigment. 

The solubility product of Pb2[Fe(CN)6] given by Krleza et al., 336 which they used as a reference to determine the solubility products, is much lower than the one used by Tananaev et al.. If applied to Tanan-aev s calculations, this produces a solubility of Iron Blue of only 0.05 mg per liter. Krleza et al., however, find similar results for the solubility of most of the metal cyanides analyzed, including Iron Blue. Since conventional methods of analysis, such as gravimetry and titration, tend to be unreliable when facing minute traces, one must but wonder about these similar results. 

Should the free Fe3+ concentration surpass this value due to a better solubility of Iron Blue, then Fe3+ would precipitate as hydroxide and would be increasingly removed from the pigment, thereby destroying it in the end. Since this does not happen at pFI=7 at all, and pH=10 can be considered the point where it just starts to happen, the concentration of the Fe3+ ion in a saturated Iron Blue solution must lie well below 10 18 mol/liter, i. e., in the area of 10"27 mol/liter. Thus, the solubility of Iron Blue must also have a value around 10 27 mol per liter (actually % of the free Fe3+ concentration, Ks less than 4.1 10 187 mol7 l 7, pKs larger than 186.6) which, at a mol mass of 1,110 g mol 1 ((Fe4[Fe(CN)6]3 14 H20) would correlate to 10 24 g. 

Several studies have investigated the effect of chelators on the Fenton reaction. Addition of chelators to Fe(III)-H202 systems allows for effective degradation at near-neutral pH values [17,23,61], Much of the influence of the iron chelators stems from increased solubility of iron species at higher pH values. These chelators were also effective for combined Fenton bioremediation treatment because they eliminate the need for low pH, which can harm useful bacteria. 

Since ascorbic acid is a strong reducing agent, it is presumed to reduce ferric iron to its ferrous form or to maintain ferrous iron in its reduced state which is more available. However, its action on copper would be to reduce Cu (III) to Cu (I), the less available form of copper (36). Ascorbic acid is also thought to Increase the solubility of iron by decreasing the alkalinity of the intestinal chyme. Effect of ascorbic acid on the availability of zinc has been less extensively studied. 

Comparison of the solubility of iron hydroxides and goethite according to the data of Lengweiler et al. (1961) and Schindler et al. (1963) showed that there is a functional relationship between the variation in isobaric potential (AAG) during aging of the precipitate and aging time (Mel nik, 1972b). This made it possible to calculate the isobaric potentials of formation of the compounds—from thermodynamically stable a-FeOOH to metastable freshly precipitated Fe(OH)3-I (AC°298, kcal/mol) ... 

Hemley et al. (1992) studied the solubility of iron, lead, zinc, and copper sulfides in chloride solutions that were rock-bufifered in pH and in oxygen and sulfur fugacity, in the range 300-700 °C 50-200 MPa. Their results show that iron-, copper-, zinc-, and lead-sulfide mineral solubilities decrease with decreasing temperature and decreasing total chloride (see Figure 2), and with increasing pressure. In nature, the HCl concentration (which is the sum of HC1° and some portion of the ionized and Cl ) in a hydrothermal solution is controlled by equilibria such as... 

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