Iron hydroxide, solubility-product

More recendy, the molten caustic leaching (MCL) process developed by TRW, Inc. has received attention (28,31,32). This process is illustrated in Eigure 6. A coal is fed to a rotary kiln to convert both the mineral matter and the sulfur into water- or acid-soluble compounds. The coal cake discharged from the kiln is washed first with water and then with dilute sulfuric acid solution countercurrendy. The efduent is treated with lime to precipitate out calcium sulfate, iron hydroxide, and sodium—iron hydroxy sulfate. The MCL process can typically produce ultraclean coal having 0.4 to 0.7% sulfur, 0.1 to 0.65% ash, and 25.5 to 14.8 MJ/kg (6100—3500 kcal/kg) from a high sulfur, ie, 4 wt % sulfur and ca 11 wt % ash, coal. The moisture content of the product coal varies from 10 to 50%. 

The extent of hydrolysis of (MY)(n 4)+ depends upon the characteristics of the metal ion, and is largely controlled by the solubility product of the metallic hydroxide and, of course, the stability constant of the complex. Thus iron(III) is precipitated as hydroxide (Ksal = 1 x 10 36) in basic solution, but nickel(II), for which the relevant solubility product is 6.5 x 10 l8, remains complexed. Clearly the use of excess EDTA will tend to reduce the effect of hydrolysis in basic solutions. It follows that for each metal ion there exists an optimum pH which will give rise to a maximum value for the apparent stability constant. 

Solutions which prevent the hydrolysis of salts of weak acids and bases. If the precipitate is a salt of weak acid and is slightly soluble it may exhibit a tendency to hydrolyse, and the soluble product of hydrolysis will be a base the wash liquid must therefore be basic. Thus Mg(NH4)P04 may hydrolyse appreciably to give the hydrogenphosphate ion HPO and hydroxide ion, and should accordingly be washed with dilute aqueous ammonia. If salts of weak bases, such as hydrated iron(III), chromium(III), or aluminium ion, are to be separated from a precipitate, e.g. silica, by washing with water, the salts may be hydrolysed and their insoluble basic salts or hydroxides may be produced together with an acid ... 

C18-0073. For the following salts, write a balanced equation showing the solubility equilibrium and write the solubility product expression for each (a) silver chloride (b) barium sulfate (c) iron(H) hydroxide and (d) calcium phosphate. 

When the post-coagulation sludge is added to hydrolyzate production, then the iron hydroxide is converted into ferric sulfate, which is well soluble in water. Too high content of soluble iron compounds, which are easily absorbable, should be avoided, as the excessive amount of the iron in the diet is harmful [13]. It leads to some disease like hemochromatosis or siderosis. Thus, the aim of our research was to find a method of decreasing the absorbable iron content of fish silage. 

Kgo is the solubility product. It applies to iron oxides, hydroxides and oxide hydroxides. 

Biedermann, G. Schindler, P. (1957) On the solubility products of precipitated iron(lll) hydroxide. Acta Chem. Scand. 11 731-740 Bigham, J.M. Ciolkosz, E.J. (1993) Soil colour. Soil Sci. Soc. Am. Spec. Publ. 31, Madison, Wl, 159 p. 

Vlek, P.L.G. Blom,Th.J.M. Beek, J. Lindsay, W.L. (1974) Determination of the solubility product of various iron hydroxides and jaro-site by the chelation method. Soil Sci Soc. Am. Proc. 38 429-432... 

It has long been recognized that ferric iron is a moderately strong acid. As early as 1896, Goodwin (5) concluded from conductometric measurements that simple dilution of ferric chloride solutions led to the formation of FeOH2+. The insolubility of ferric hydroxide has of course been appreciated even longer. The best current estimate of the solubility product constant for Fe OH)s at 25° (in 3 M NaC104) is (d). 

Metastability of Hydrolyzed Iron (III) Solutions The low solubility of ferric hydroxide has been alluded to in the Introduction. Feitknecht and Michaelis (29) have observed that aU ferric perchlorate solutions to which base has been added are unstable with respect to eventual precipitation of various forms of hydrated ferric oxides. In 3 M NaC104 at 25° C the two phase system reaches an apparent equilibrium after 200 hours, according to Biedermann and Schindler (6), who obtained a reproducible solubility product constant for ferric hydroxide at varying degrees of hydrolysis. It appears that many of the solutions used in the equilibrium studies of Hedstrom (9) and Biedermann (22) were metastable, and should eventually have produced precipitates. Nevertheless, since the measured potentials were reversible, the conclusions reached about the species present in solution remain valid. 

Because electrode potentials are defined with reference to the H+/H2 electrode under standard conditions, E° values apply implicitly to (hypothetically ideal) acidic solutions in which the hydrogen ion concentration is 1 mol kg-1. Such E° values are therefore tabulated in Appendix D under the heading Acidic Solutions. Appendix D also lists electrode potentials for basic solutions, meaning solutions in which the hydroxide ion concentration is 1.0 mol kg-1. The conversion of E° values to those appropriate for basic solutions is effected with the Nernst equation (Eq. 15.15), in which the hydrogen ion concentration (if it appears) is set to 1.0 x 10-14 mol kg-1 and the identity and concentrations of other solute species are adjusted for pH 14. For example, for the Fc3+/2+ couple in a basic medium, the relevant forms of iron(III) and iron(II) are the solid hydroxides, and the concentrations of Fe3+ (aq) and Fe2+ (aq) to be inserted into the Nernst equation are those determined for pH 14 by the solubility products of Fe(OH)3(s) and Fe(OH)2(s), respectively. Examples of calculations of electrode potentials for nonstandard pH values are given in Sections 15.2 and 15.3. 

A major problem for cells is posed by the relative insolubility of ferric hydroxide and other compounds from which iron must be extracted by the organism. A consequence is that iron is often taken up in a chelated form and is transferred from one organic ligand, often a protein, to another with little or no existence as free Fe3+ or Fe2+. As can be calculated from the estimated solubility product of Fe(OH)3 (Eq. 16-1),7 the equilibrium concentration of Fe3+ at pH 7 is only 10-17 M. 

Eventually this process forms the neutral species Fe(H2O)3(OH)30, which precipitates as amorphous iron hydroxide, which may settles out of the water column. Figure 3 shows the predicted effect of pH on the relative concentrations of the various iron hydrolysis species with and without considering the iron hydroxide solid, which dominates the speciation above pH 3.0 at 1 dM total iron. The log of the solubility product of this solid is -38.8, indicating that iron is very insoluble at natural pH values. Over time, this metastable amorphous material converts to more thermodynamically... 

Ammonia solution precipitation of iron(II) hydroxide occurs (cf. reaction 1). If, however, larger amounts of ammonium ions are present, the dissociation of ammonium hydroxide is suppressed, (cf. Section 1.15), and the concentration of hydroxyl ions is lowered to such an extent that the solubility product of iron(II) hydroxide, Fe(OH)2, is not attained and precipitation does not occur. Similar remarks apply to the other divalent elements of Group III, nickel, cobalt, zinc and manganese and also to magnesium. 

A suite of both oxidized and reduced iron minerals has been found as efflorescences and precipitates in or near the acid mine water of Iron Mountain. The dominant minerals tend to be melan-terite (or one of its dehydration products), copiapite, jarosite and iron hydroxide. These minerals and their chemical formulae are listed in Table III from the most ferrous-rich at the top to the most ferric-rich at the bottom. These minerals were collected in air-tight containers and identified by X-ray diffractometry. It was also possible to check the mineral saturation indices (log Q(AP/K), where AP = activity product and K = solubility product constant)of the mine waters with the field occurrences of the same minerals. By continual checking of the saturation index (S.I.) with actual mineralogic occurrences, inaccuracies in chemical models such as WATEQ2 can be discovered, evaluated and corrected (19), provided that these occurrences can be assumed to be an approach towards equilibrium. 

Iron of inorganic dissolved compounds (bicarbonates, sulfates, chlorides, fluosilicates, etc.) may enter into the dissolved form of iron of inorganic origin (Fcj"" ), but their existence is governed by an acid environment with a pH not higher than 3. As a rule the pH in sea water is close to 8 ( 0.5). Under these conditions iron compounds are easily hydrolyzed and converted into hydroxides, which form colloidal solutions in sea water. In appropriate conditions colloidal hydroxide condenses to clots of gel and converts to the suspended state. Therefore there are practically no ionic forms of iron (Fe "" proper). As early as 1937 Cooper (1937) concluded, on the basis of the solubility product and activity of ferrous and ferric iron ions and FeOH ions, that until equilibrium is reached sea water may contain about 10 jiig/1 of iron ions in true solution at pH = 8.5 the amount of ionic Fe in ferric form is still less—10 which corresponds to the extremely... 

Schellmann (unpublished results) calculated from the solubility products of manganese and iron hydroxides that interstitial water can be sufficiently supplied with manganese and iron even at higher redox potentials, because the time of dissolution is adequately long. 

Loughnan (1969) discussed the solubility in relation to pH of some of the common products of chemical weathering of silicate minerals, In general, the hydroxides of Na, K and Ca are soluble at all pH s, and Mg(OH)2 is soluble at pH < 10. Aluminium oxide is soluble at pH s < 4 and >10, whereeis SiOj is slightly soluble at pH < 9 and increasingly soluble at higher pH values. Titanium hydroxide is soluble at pH < 5, but TiOz is soluble only at pH < 2. The hydroxide of trivalent iron is soluble only below pH 2.5, but Fe(OH)2 is soluble below about pH 8.5. 

The formation of a monovalent species in deposition and dissolution of divalent ions of the iron-group transition metals is commonly assumed in the literature. Since the monovalent ions (such as Fe ) are unstable in solution, they are assumed to be adsorbed on the surface, either as the ion itself or as a hydroxide, such as FeOHads- This could stabilize the monovalent form of the element. Moreover, since a monovalent hydroxide is not known to exist in solution, one does not know its solubility product, and the possibility of the existence of this adsorbed species, even in solutions of low pH, cannot be ignored a priori. Unfortunately, the nature of the adsorbed intermediate, or even the evidence for its existence on the surface, is at best circumstantial. 

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