Solubility of metal hydroxides

FIGURE 6.3 Solubility of metal hydroxides and sulfides. (Taken from Krofta, M. and Wang, L.K., Design of Innovative Flotation-Filtration Wastewater Treatment Systems for a Nickel-Chromium Plating Plant, U.S. Department of Commerce, National Technical Information Service, Springfield, VA, Technical Report PB-88-200522/AS, January 1984.)... 

Figure 7 Solubilities of metal hydroxides and metal sulfides. (Courtesy of USEPA.)...

Solubility of Metal Hydroxides and Sulfides as a Function of pH. (Source EPA publication, EPA-600/2-82-011C, 1981)... 

The solubility of metal-hydroxide precipitates in water varies depending on ionic strength and number of pairs and/or complexes (Chapter 2). A practical approach to determining the pH of minimum metal-hydroxide solubility, in simple or complex solutions, is potentiometric titration, as demonstrated in Figure 12.3. The data show that potentiometric titration of a solution with a given heavy metal is represented by a sigmoidal plot. The long pH plateau represents pH values at which metals precipitate the equivalence point, or titration end point, indicates the pH at the lowest metal-... 

Solubility of metal hydroxides increases at lower pH because the acid present ties up hydroxide ions in solution and so shifts the dissolution equilibrium to the right. 

While ranges of total concentration serve to set bounds for experimentally determining effects on marine populations, the actual species of metal ion available to the biological population is of importance. Sillen, in a classic paper, has computed the stable species of many metals in sea water21). He concluded, for example, that Hg+2, Cd+2, and Pb+2 exist primarily as chloride complexes. pH determines the availability of the hydroxide ion and thereby the solubility of metal hydroxides. Sillen assumed a pH of 8.1 0.2 as representative. Significant variations could occur, however, in estuarine waters. When concentrations of trace elements were compared with calculations of their solubility products and stability constants, the observed values were considerably less than the calculated values. The implication is that the heavy metals are not in equilibrium with solid phases of their salts, but that other processes, such as chelation and adsorption, control their concentration. 

The hydroxides M (OH)2 are generally less soluble and are of lower base strength. The Group I hydroxides are almost unique in possessing good solubility—most metal hydroxides are insoluble or sparingly soluble hence sodium hydroxide and, to a lesser extent potassium hydroxide, are widely used as sources of the hydroxide ion OH" both in the laboratory and on a large scale. 

Values for the solubility products of metallic hydroxides are, however, not very precise, so that it is not always possible to make exact theoretical calculations. The approximate pH values at which various hydroxides begin to precipitate from dilute solution are collected in Table 11.2. 

Most of the pollutants may be effectively removed by precipitation of metal hydroxides or carbonates using a reaction with lime, sodium hydroxide, or sodium carbonate. For some, improved removals are provided by the use of sodium sulfide or ferrous sulfide to precipitate the pollutants as sulfide compounds with very low solubilities. After soluble metals are precipitated as insoluble floes, one of the water-solid separators (such as dissolved air flotation, sedimentation, centrifugation, membrane filtration, and so on) can be used for floes removal.911 The effectiveness of pollutant removal by several different precipitation methods is summarized in Tables 5.15-5.17. 

These changes have significant effects on the solubility of metals in the KF solution. Firstly, the reduction in pE caused some reduction of Fe(OH)3 to Fe", increasing the Fe concentration in the digestion and lowering the stability boundary between Fe" and Fe(OH)3. Secondly, the higher pHs resulted in less adsorption of Pb and Zn, and possibly the precipitation of Pb and Zn hydroxides, resulting in less Pb and Zn in solution and more concentrations below detection. 

The coordination atmosphere of the metal ion in solution can also be expected to affect the reaction rate. Microanalytical results indicate that the active catalysts in cobalt and nickel systems could well be metal thiolic species produced in situ. However, these complexes are appreciably more soluble in the, alkaline solutions than are metal hydroxides (see, for example, the analysis results reported in Table IV), and it is not possible on the present evidence to differentiate between catalysis as a result of increased solubility (comparing metal hydroxides and metal thiolic complexes), and catalysis as a result of differences in the allowed ease of electron transfer. It is apparent, however, that most of the metals investigated (Table I) are poor catalysts because they form only the insoluble hydroxide complexes. 

Dosing with alkali (usually hydrated lime (Ca(OH)2), and less frequently caustic soda (NaOH)), which serve both to raise the pH (thus lowering the solubility of most problematic metals) and to supply OH ions for the rapid precipitation of metal hydroxide solids. 

The exceptional character o fluorine.—Fluorine has a little more individuality, so to speak, than the other three members of the family (1) There are no compounds of oxygen and fluorine (2) Chlorine, bromine, and iodine or the haloid acids show no signs of the remarkable effect of hydrofluoric acid and of fluorine on silicon (3) The solubilities of the sulphates, nitrates, and chlorides of barium, strontium, calcium, and magnesium decrease with increasing at. wt. of the metal, while the solubilities of the hydroxides increase the solubilities of the iodides, bromides, and chlorides... 

Further evidence has been obtained to support the contention that the active catalysts are metal complexes dissolved in solution. With experiments reported in Table II, the kinetics of oxidation under standard conditions in the presence of various metal salts are compared with the rates of reaction when solid residues have been filtered from solution. The agreement between the rates in Cases 1 and 3 of Table II (where the amount of metal available is dictated by the solubility of metal complexes) shows that solid precipitates play little or no part in catalysis in all the systems studied. The amount of metal in solution has been measured in Cases 2 and 3 metal hydroxide complexes (Case 2) are not as soluble as metal-thiol complexes, and neither is as soluble as metal phthalocyanines (19). The results of experiments involving metal pyrophosphates are particularly interesting, in that it has previously been suggested that cobalt pyrophosphates act as heterogeneous catalysts. The result s in Table II show that this is not true in the present system. 

Precipitation of Metal Hydroxides. In general the addition of a soluble base to the solution of a metal salt produces a precipitate. This precipitate is of rather variable composition, but its nature is best understood if it is regarded as the hydroxide of the metal. Thus, sodium hydroxide added to copper sulphate solution gives a light blue voluminous precipitate and sodium hydrox-... 

Magnesium has a very strong affinity for oxygen as shown by the intensity with which the metal burns. However, the residual affinity of magnesium oxide for water is much smaller than that of calcium oxide. The solubility of magnesium hydroxide is so small that the saturated solution acts but slowly in turning litmus blue. 

The pH-dependent solubility behavior of metal-hydroxides and the corresponding solution pH of zero net charge can be demonstrated by deriving pH dependent solubility functions for all the metal-hydroxy species of a particular metal in solution. 

Traditionally, the objective of precipitation is to obtain a well-dispersed metal hydroxide (or carbonate) phase on a support through its precipitation from an aqueous metal solution onto a support powder by adding a base. The support powder is suspended in the metal salt solution. Upon addition of, say, NaOH, the pH increases strongly and metal hydroxide species are generated. When their concentration exceeds the (super)solubility limit, metal hydroxide particles can... 

Precipitation and dissolution of metal hydroxides The solubility product principle can also be applied to the formation of metal hydroxide precipitates these are also made use of in qualitative inorganic analysis. Precipitates will be formed only if the concentrations of the metal and hydroxyl ions are momentarily higher than those permitted by the solubility product. As the metal-ion concentration in actual samples does not vary much (10—1 —10 3 mol -1 is the usual range), it is the hydroxyl-ion concentration which has the decisive role in the formation of such precipitates. Because of the fact that in aqueous solutions the product of hydrogen- and hydroxyl-ion concentrations is strictly constant (A = 10 14 at 25°C, cf. Section 1.18), the formation of a metal-hydroxide precipitate depends mainly on the pH of the solution. Using the solubility product principle, it is possible to calculate the (minimum) pH required for the precipitation of a metal hydroxide. 

It is useful to state that there are complete separation systems of cations devised, based entirely on the separation of metal hydroxides, omitting the use of sulphides altogether. These however will not be discussed here. As it will be seen later, the separation of ions within the third group of cations, is almost entirely based on the differences in the solubilities of their hydroxides. 

The precipitation method is not only applicable to one hydroxide in ca.se a solution of mixed salts is used, coprecipitation takes place [54]. As in conventional coprecipitation, the solubilities of various hydroxides may not be the same, leading to a molecular ratio of the two metals in the deposited layer being different from that in the solution. Due to difficulties with respect to stirring, concentration gradients may exist that will in turn influence the homogeneity of the precipitated phase. The pore volume of the substrate may restrict the amount deposited. 

In the Me -H20-C02 system in the presence of the earth s atmosphere, carbonates are frequently more stable than oxides or hydroxides as solid phases. Thus in natural water systems the concentration of some metal ions is controlled by the solubility of metal carbonates. 

Iron, another constituent of cement, was present in the extracts at concentrations which followed the solubility of ferric hydroxide at the pH of the leachate. Iron is known to be much less soluble than aluminum in this pH range, and this was reflected in the lower amounts in solution. Iron behaved similarly to the less soluble heavy metals, chromium and lead. 

In other cases the conductivities of redissolved hydroxide solutions are equal to those of alkali at the same concentration, suggesting that the solubility of the hydroxide is more probably due to the formation of a sol. For example, the freshly precipitated cream-coloured Ge(OH)j dissolves in strongly basic solutions to form reddish-brown colloidal solutions. ) Hydrolysis of metal-salt solutions often yields colloidal solutions of the hydroxides, as in the case of the trihydroxides of Fe and Cr and the tetrahydroxides of elements of the fourth Periodic Group such as Sn, Ti, Zr, and Th. It is unlikely that the hydroxides M(OH)4 of the latter elements are present in such solutions, for the water content of gels MO2. XH2O is very variable. 

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