Complex-ion formation constants

In common with other hydroxy organic acids, tartaric acid complexes many metal ions. Formation constants for tartaric acid chelates with various metal ions are as follows Ca, 2.9 Cu, 3.2 Mg, 1.4 and Zn, 2.7 (68). In aqueous solution, tartaric acid can be mildly corrosive toward carbon steels, but under normal conditions it is noncorrosive to stainless steels (Table 9) (27). 

Bromine is moderately soluble in water, 33.6 g/L at 25°C. It gives a crystalline hydrate having a formula of Br2 <7.9H2 O (6). The solubiUties of bromine in water at several temperatures are given in Table 2. Aqueous bromine solubiUty increases in the presence of bromides or chlorides because of complex ion formation. This increase in the presence of bromides is illustrated in Figure 1. Kquilibrium constants for the formation of the tribromide and pentabromide ions at 25°C have been reported (11). 

Formation constants are the equilibrium constants for complex ion formation. The stepwise formation constants, designated K, are defined as follows ... 

From Eqn. (14) it follows that with an exothermic reaction - and this is the case for most reactions in reactive absorption processes - decreases with increasing temperature. The electrolyte solution chemistry involves a variety of chemical reactions in the liquid phase, for example, complete dissociation of strong electrolytes, partial dissociation of weak electrolytes, reactions among ionic species, and complex ion formation. These reactions occur very rapidly, and hence, chemical equilibrium conditions are often assumed. Therefore, for electrolyte systems, chemical equilibrium calculations are of special importance. Concentration or activity-based reaction equilibrium constants as functions of temperature can be found in the literature [50]. 

Consider complex ion formation in the CdClj-KCl system, and let it be assumed for the moment that a CdCl complex ion is formed. If such complex ions were formed in an aqueous solution of CdClj and KCl, they would exist as little islands separated from other ions by large expanses of water. In fused salts, there are no oceans of solvent separating the ions. Thus, a Cd " ion would constantly be coming into contact on all sides with chloride ions, and yet one singles out three of these CP ions and says that they are part of (or belong to) a CdCIJ complex ion (Fig. 5.54). It appears that in the absence of the separateness possible in aqueous solutions, the concept of complex ions in molten salts is suspect As will be argued later, however, what is dubious turns out to be not the concept but the comparison of complex formation in fused salts with complex formation in aqueous solutions. 

A measure of the tendency of a metal ion to form a particular complex ion is given by the formation constant Kf (also called the stability constant), which is the equilibrium constant for the complex ion formation. The larger Kf is, the more stable the complex ion is. Table 16.4 lists the formation constants of a number of complex ions. 

Complex ions are formed in solution by the combination of a metal cation with a Lewis base. The formation constant Kf measures the tendency toward the formation of a specific complex ion. Complex ion formation can increase the solubility of an insoluble substance. 

Formation constant. The equilibrium constant for the complex ion formation. (16.10)... 

To see the effect of complex-ion formation on the solubility of ZnS, we add the equations and, therefore, multiply their equilibrium constants ... 

The complexing of zirconium and hafnium ions by fluoride ions is quite extensive compared to chloro complexing, while complex ion formation w ith bromide and iodide ions is negligible. Formation constants for fluoride complexing with zirconium(IV) and hafnium(IV) calculated from the data of Connick (126), Buslaev (94), and Hume (574), have been summarized graphically by Goldstein (213). Slightly different values have been published by Bukhsh (92). Noren (15a, 401-403) has redetermined the equilibrium constants for the reaction. 

The application of mineral solubility data (Figure 4) to soils and sediments is at best semi-quantitative because of unknown eflFects of complex ion formation, solid solution, and armoring. The apparent solubility of Fe(OH)3(a) has been reported to be essentially constant at about 10" M between a pH of 6 and 11, presumably because of undissociated Fe(OH).3 (70). Armoring by iron and aluminum oxides was believed to be responsible for the reduced solubilization rate of rounded quartz... 

So pronounced is the chelating tendency of the diketonate anion that even alkali metal complexes may be isolated, as illustrated by RbiKCFjCOCHCOCFjljNa], in which sodium is surrounded by a trigonal prismatic array of donor oxygen atoms. For complexes derived from dibenzoylmetbane, stabilities in dioxane-water are in the order Li > Na > K > Cs. For divalent metal ions formation constants increase in the order Ba > Sr > Ca > Mg > Cd > Mn > Pb > Zn > Co, Ni, Fe > Cu, and for higher valent metal ions the first formation constants for chelates are in the order Fe + > Ga " " > Th" > In " " > Sc " " > Y " " > Sm " > Nd " > More recently, such stu-... 

Some polystyrene resins (cross-linked with DVB) are specially modified to have chelating functional groups bound to the matrix so as to make them selective towards certain ions. Such resins with iminodiacetic acid groups are marketed under the trade names Dowex A-1 (Dow Chemical) and Chelex 100 (Bio-Rad Laboratories). The complex (XXVI) formation constants with metal ions of the chelating resin are so large that the resin absorbs metal ions equivalent to the iminodiacetic acid groups (used in sodium salt form), i.e., the efficiency of metal ion adsorption is near 100%. A particular metal ion can be removed by controlling the pH of aqueous solution. For example, at pH 2, mercury and copper ions are... 

The solubility of metal salts is also affected by the presence of certain Lewis bases that react with metal ions to form stable complex ions. Complex-ion formation in aqueous solution involves the displacement by Lewis bases (such as NH3 and CN ) of water molecules attached to the metal ion. The extent to which such complex formation occurs is expressed quantitatively by the formation constant for the complex ion. Amphoteric oxides and hydroxides are those that are only slightly soluble in water but dissolve on addition of either acid or hase. 

Another difference lies in the polyfimctional nature of metal ions. They are capable of reacting with more than one ligand, forming a series of complexes whose equilibrium constants are much more closely spaced than those of most polyprotic acids. For example (Figure 9.1), the titration of Cu with NH3 would certainly be of dubious practical analytical value. Copper(II) ion forms a series of ammonia complexes whose formation constants are both small and closely spaced there are several complex species present in significant concentrations throughout the titration. 

Because of the ability of ion-selective electrodes to measure the activity of free ions, it has been possible to obtain thermodynamic formation constants and to have information on the number of binding sites and the stoichiometries of complexes. Thus, formation constants have been calculated for copper(II) complexes with glycine, glutamic acid and tris(hydroxymethyl)aminomethane [420,421]. 

The formation constant for Ag(NH3)J is quite large, so most of the silver ions exist in the complexed form. In the absence of anunonia we have, at equilibrium, [Ag ] = [Cr]. As a result of complex ion formation, however, we can write... 

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