Solubility equilibria with complex ions

The chemical species present in the electrolyte are actually more complex than that described. In solution, elemental bromine exists in equilibrium with bromide ions to form polybromide ions, Br, where = 3, 5, 7. Aqueous zinc bromide is ionized, and zinc ions exist as various complex ions and ion pairs. The electrolyte also contains complexing agents which associate with polybromide ions to form a low-solubility second liquid phase. The complex reduces the amount of bromine contained in the aqueous phase 10 to 100-fold, which, in addition to the separator, also reduces the amount of bromine available in the eeU for the self-discharge reaction. The complex also provides a way to store bromine at a site remote from the zinc deposits and is discussed further in the next section. Salts with organic cations such as iV-methyl-iV-ethylmorpholinium bromide (MEMBr) are commonly used as the complexing agents. One researcher has proposed a mixture of four quaternary ammonium salts for use in zinc/bromine batteries. The proposed electrolyte has favorable properties with regard to aqueous bromine concentration, resistivity, and bromine diffusion and does not form solid complexes at low temperatures (5°C and above). Complexes with quaternary ammonium ions are reversible and also have an added safety benefit due to a much reduced bromine vapor pressure (see Sec. 35.6). 

BrCl exists in equilibrium with bromine and chlorine in both gas and liquid phases. Table 5 lists various physical properties of BrCl. Due to the polarity of BrCl, it shows greater solubility than bromine in polar solvents. In water, it has a solubility of 8.5 gms per 100 gms of water at 20 C (that is, 2.5 times the solubility of bromine 11 times that of chlorine). Bromine chloride s solubility in water is increased greatly by adding chloride ions to form the complex chlorobromate ion, BrCl 2. 

K has the value of about 1 x 10 at 298 K, and in solutions of copper ions in equilibrium with metallic copper, cupric ions therefore greatly predominate (except in very dilute solutions) over cuprous ions. Cupric ions are therefore normally stable and become unstable only when the cuprous ion concentration is very low. A very low concentration of cuprous ions may be produced, in the presence of a suitable anion, by the formation of either an insoluble cuprous salt or a very stable complex cuprous ion. Cuprous salts can therefore exist in contact with water only if they are very sparingly soluble (e.g. cuprous chloride) or are combined in a complex, e.g. [Cu(CN)2) , Cu(NH3)2l. Cuprous sulphate can be prepared in non-aqueous conditions, but because it is not sparingly soluble in water it is immediately decomposed by water to copper and cupric sulphate. 

The complexation of Pu(IV) with carbonate ions is investigated by solubility measurements of 238Pu02 in neutral to alkaline solutions containing sodium carbonate and bicarbonate. The total concentration of carbonate ions and pH are varied at the constant ionic strength (I = 1.0), in which the initial pH values are adjusted by altering the ratio of carbonate to bicarbonate ions. The oxidation state of dissolved species in equilibrium solutions are determined by absorption spectrophotometry and differential pulse polarography. The most stable oxidation state of Pu in carbonate solutions is found to be Pu(IV), which is present as hydroxocarbonate or carbonate species. The formation constants of these complexes are calculated on the basis of solubility data which are determined to be a function of two variable parameters the carbonate concentration and pH. The hydrolysis reactions of Pu(IV) in the present experimental system assessed by using the literature data are taken into account for calculation of the carbonate complexation. 

Hydrogen cyanide (Table 15.1) is a colorless, flammable liquid or gas that boils at 25.7°C and freezes at minus 13.2°C. The gas rarely occurs in nature, is lighter than air, and diffuses rapidly. It is usually prepared commercially from ammonia and methane at elevated temperatures with a platinum catalyst. It is miscible with water and alcohol, but is only slightly soluble in ether. In water, HCN is a weak acid with the ratio of HCN to CN about 100 at pH 7.2, 10 at pH 8.2, and 1 at pH 9.2. HCN can dissociate into H+ and CN. Cyanide ion, or free cyanide ion, refers to the anion CN derived from hydrocyanic acid in solution, in equilibrium with simple or complexed cyanide molecules. Cyanide ions resemble halide ions in several ways and are sometimes referred to as pseudohalide ions. For example, silver cyanide is almost insoluble in water, as are silver halides. Cyanide ions also form stable complexes with many metals. 

Speciation is a dynamic process that depends not only on the ligand-metal concentration but on the properties of the aqueous solution in chemical equilibrium with the surrounding solid phase. As a consequence, the estimation of aqueous speciation of contaminant metals should take into account the ion association, pH, redox status, formation-dissolution of the solid phase, adsorption, and ion-exchange reactions. From the environmental point of view, a complexed metal in the subsurface behaves differently than the original compound, in terms of its solubility, retention, persistence, and transport. In general, a complexed metal is more soluble in a water solution, less retained on the solid phase, and more easily transported through the porous medium. 

Almost all analyte ion inside the membrane in Figure 15-8b is bound in the complex LC+, which is in equilibrium with a small amount of free C+ in the membrane. The membrane also contains excess free L. C+ can diffuse across the interface. In an ideal electrode, R cannot leave the membrane, because it is not soluble in water, and the aqueous anion A-cannot enter the membrane, because it is not soluble in the organic phase. As soon as a few C+ ions diffuse from the membrane into the aqueous phase, there is excess positive charge in the aqueous phase. This imbalance creates an electric potential difference that opposes diffusion of more C+ into the aqueous phase. 

If there is introduced into the solution from some other source an ion that is in common with an ion of the insoluble solid, the chemical equilibrium is shifted to the left, and the solubility of that solid will be greatly decreased from what it is in pure water. This is called the 11 common-ion effect." This effect is important in gravimetric analysis, where one wishes to precipitate essentially all of the ion being analyzed for, by adding an excess of the "common-ion" precipitating reagent. There is a practical limit to the excess, however, which involves such factors as purity of precipitate and possibility of complex formation. You can calculate the solubility under a variety of conditions, as illustrated in the following problem. 

Two simultaneous equilibria are involved the solubility equilibrium (horizontal equation), and the dissociation of the complex ion (vertical equation). The Ag+ is shared in common with both equilibria. 

The solubility of an ionic compound increases dramatically if the solution contains a Lewis base that can form a coordinate covalent bond (Section 7.5) to the metal cation. Silver chloride, for example, is insoluble in water and in acid, but it dissolves in an excess of aqueous ammonia, forming the complex ion Ag(NH3)2 + (Figure 16.13). A complex ion is an ion that contains a metal cation bonded to one or more small molecules or ions, such as NH3, CN-, or OH-. In accord with Le Chatelier s principle, ammonia shifts the solubility equilibrium to the right by tying up the Ag+ ion in the form of the complex ion ... 

A similar consideration can be applied to the cathodic processes. In a solution of mercuric nitrate bivalent mercury will bo reduced to univalent until the ratio of the respective activity of the mercurous salt formed and tho mercuric salt still remaining reaches the equilibrium value. During the course of further reaction the ratio of activities of both ions in the solution will not change any longer, and metallic mercury will be deposited. Therefrom, it is evident that mercuric nitrate cannot be quantitatively reduced to mercurous salt. Bivalent mercury can be reduced practically completely to univalent in the case of mercuric chloride. As the solubility of the mercurous chloride formed by the reduction and consequently also the concentration of Hg2+ ion is very small the equilibrium between the ions in the solution will be attained only then, when nearly all Hg++ ions will be reduced to univalent ones. On the other hand on reduction of the very slightly dissociated cyanide complex Hg(CN) the equilibrium between mercurous and mercuric ions is reached at the very beginning of electrolysis as soon as a hardly noticeable amount of Hg++ ions has been formed from that moment on metallic mercury will be deposited at the cathode with practically 100 p. o. yield. 

How can we pull the dissolution equilibrium to the right, even though CP is an extremely weak base The key is to lower the concentration of Ag+ in solution by forming complex ions. For example, Ag+ reacts with excess NH3 to form the stable complex ion Ag(NH3)2+. As a result, AgCl is quite soluble in concentrated ammonia solutions. The relevant reactions are... 

However, the case in which the solubility of a solid can be calculated from the known analytical concentration of added components and from the solubility product alone is very seldom encountered. Ions that have dissolved from a crystalline lattice frequently undergo chemical reactions in solution, and therefore other equilibria in addition to the solubility product have to be considered. The reaction of the salt cation or anion with water to undergo acid-base reactions is very common. Furthermore, complex formation of salt cation and salt anion with each other and with one of the constituents of the solution has to be considered. For example, the solubility of FeS(s) in a sulfide-containing aqueous solution depends on, in addition to the solubility equilibrium, acid-base equilibria of the cation (e.g., Fe + H2O = FeOH + H ) and of the anion (e.g., S + HjO = HS + OH, and HS" + H2O = H2S + OH ), as well as on equilibria describing complex formation (e.g., formation of FeHS" or FeSi ). 

The relations depicted in Figure 7.3 or characterized by equation 14 do not fully describe the solubility of oxides or hydroxides. We have to consider that the solid can be in equilibrium with hydroxo-metal-ion complexes [Me(OH)J -". 

Figure 7.3. Solubility of oxides and hydroxides. Free metal-ion concentration in equilibrium with solid oxides or hydroxides. The occurrence of hydroxo metal complexes must be considered for evaluation of complete solubility.

The procedure outlined above is fairly typical for a number of determinations of solubility products of metal selenites. Data from such investigations have been reevaluated by the review with the accepted protonation constants of the selenite ion, corrected for the hydrolysis of the metal ion when necessary, and the value of the solubility product extrapolated to standard state conditions. It has been observed that the initial and final pH values in cases are in conflict. This has been ignored and the calculations have been based solely on the data for the equilibrium solution. Complexation of the metal ion by the anions and temperature effects were neglected, which probably introduces a negligible error compared with other sources of error. Activity coefficients were calculated by the SIT expression with s = 0 kg-mol, which is a reasonable simplification due to the low ionic concentrations. The ionic strength was obtained by an iterative procedure from knowledge of the total metal concentration and the pH of the equilibrium solution. The results of the recalculations are entered in Chapter V. 

Both sets of compounds are soluble in water, but solutions of Cd halides contain a variety of complex ions [CdXJ in equilibrium. Hg halides have varying coordination, with two close... 

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 equilibrium process can be influenced by several factors which include adjustment of pH to prevent ionization of acids or bases, by the formation of ion-pairs with ionizable analytes, by the formation of hydrophobic complexes with metal ions, or by adding neutral salts to the aqueous phase to reduce the solubility of the analyte (also known as salting out ). 

In our spreadsheet exercise we will consider a textbook example, the solubility of HgS as a function of pH, in a solution in equilibrium with solid HgS that contains no other sources of mercury and sulfur. HgS is quite insoluble, with a reported solubility product of 5 X 10 54 M2. The case is complicated by the fact that the two participant ions, Hg2+and S2 , are both involved in acid-base equilibria. For Hg2+these are the successive formation of three hydroxy complexes HgOH+, Hg(OH)2, and Hg(OH)3 , for S2 the consecutive... 

Next, consider the suggestion that copper corrodes in the concentrated HC1 because of the formation of a soluble chloride complex with an equilibrium constant for the reaction Cu2+ + 4CL = (CuCl4)2- of K = 10+6. If a CuC1 2- = KL4, and the activity of the CL is that given above in the concentrated acid (acr = 5), calculate Ecell and determine whether corrosion will occur due to the formation of the complex ion. Cell reaction ... 

The manganate ion is not reduced by bromide ion but is reduced slowly by iodide ion and quickly by vanadyl(IV) or hexacyano-ferrate(II) ions. When the latter two ions are used as reductants, especially with the potassium complex, green products are obtained rapidly and in high yield. The green species is unstable in solution and is apparently in equilibrium with the reactants. With potassium salts, the solubility of the product is low, and the reaction is driven to completion. Potentiometric titrations show that a one-electron reduction occurs to produce the green species, which has been characterized by analysis and optical and e.s.r. spectroscopy. It is a mixed-valence species similar to the heteropoly blues of molybdenum and tungsten. E.s.r. spectra suggest that the extra electron is fairly well trapped on a specific vanadium atom, and the complex is therefore a class II mixed-valence species.8... 

The strength of association between the ions in solution is expressed by various equilibrium constants. Stability (formation) constants refer to complex ions and ion pairs hydrolysis (deprotonation) constants refer to the loss of H+ from the water ligands surrounding central cations. Solubility products refer to the aqueous ion activities in equilibrium with solid phases. Some constants are reported in the literature in terms of concentrations rather than activities. Such constants are misnamed, since they depend both on the concentration and on the nature of other ions in solution. Converting concentrations to activities gives a much more useful value. 

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