Ionic equilibria complex ions

Table 55 presents the results discussed above. Fluoride melts containing tantalum contain two types of complex ions, namely TaF6 and TaF72 . The equilibrium between the complexes depends on the concentration of fluoride ions in the system, but mostly upon the nature of the outer-sphere cations. The complex ionic structure of the melts can be adjusted by adding cations with a certain polarization potential. For instance, the presence of low polarization potential cations, such as cesium, leads primarily to the formation of TaF72 complexes, while the addition of cations with relatively high polarization potentials, such as lithium or sodium, shifts the equilibrium towards the formation of TaF6 ions. 

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 ... 

The solubility product, Ksp, for an ionic compound is the equilibrium constant for dissolution of the compound in water. The solubility of the compound and Ksp are related by the equilibrium equation for the dissolution reaction. The solubility of an ionic compound is (1) suppressed by the presence of a common ion in the solution (2) increased by decreasing the pH if the compound contains a basic anion, such as OH-, S2-, or CO32- and (3) increased by the presence of a Lewis base, such as NH3, CN-, or OH-, that can bond to the metal cation to form a complex ion. The stability of a complex ion is measured by its formation constant, Kf. 

Acid dissociation constants and dissociation constants of complex ions determine the concentrations of species that are present in a solution at equilibrium under specified conditions. Ionic dissociation reactions occur rapidly and tend to remain at equilibrium during an enzyme-catalyzed reaction. Since ATP (see Fig. 1.1) is the primary carrier of energy in biochemical systems and since a good deal is known about its binding properties, these properties are considered here in some detail. 

In this chapter we have seen that acid dissociation constants are needed to calculate the dependence of apparent equilibrium constants on pH. In Chapter 3 we will discuss the calculation of the effects of ionic strength and temperature on acid dissociation constants. The database described later can be used to calculate pKs of reactants at 298.15 K at desired ionic strengths. Because of the importance of pKs of weak acids, Table 1.3 is provided here. More experimental measurements of acid dissociation constants and dissociation constants of complex ions with metal ions are needed because they are essential for the interpretation of experimental equilibrium constants and heats of reactions. A major database of acid dissociation constants and dissociation constants of metal ion complexes is provided by Martell, Smith, and Motekaitis (2001). 

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]. 

Spectroscopic studies also provide evidence of the presence of some complex ions in such melts [274-277], Results obtained by Raman spectroscopy showed the presence of a residual ionic lattice composed of polynuclear aggregates of the formula (MgCl2) . Recent Raman spectroscopic investigations by Brooker et al. [278-280] showed that MgCl42" is the predominant species in MgCl2-rich melts, while other authors [281] found that the presence of small amounts of other species such as MgCl3" cannot be excluded. An equilibrium reaction,... 

We make a distinction between two types of ionic pentacoordinate complexes those which are in dynamic equilibrium with neutral hexacoordinate complexes have been dealt with in Sections III.A.4, III.A.5.ii, and III.B.2. The second group includes those pentacoordinate siliconium-ion salts which are formed as such and are stable and do not equilibrate (to a noticeable extent) with their hypothetical neutral hexacoordinate counterparts. The present section discusses this group of persistent salts of pentacoordinate silicon cations. 

The idea that the instantaneous equilibrium between the different conformations of dissymmetric complex ions could be influenced by ion-pair formation finds some support in recent experiments of Mason and coworkers with diastereoisomeric cobalt(III) complexes. Thus, Mason and Norman (31) have shown that the circular dichroism spectra of D-Co(d-pn)3" and L-Co(d-pn)3" (pn = propylenediamine) are differently changed by oxo-anions such as P04, S04 and 8203 . According to these authors the oxo-anion and D-Co(d-pn)3+ (or D-Coen3" ), but not L-Co(d-pn)3 should have a preferred mutual orientation in the ion pairs. At low concentrations the effect of sulfate and thiosulfate ions on the circular dichroism of D-Co(d-pn)3" is similar to that of phosphate but with these anions at concentrations > 0.2M the previous changes are reversed— probably due to breaking down the specific orientation of the ion pair in the denser ionic atmosphere. It is also interesting that Mason and Norman (32) have found that Co(NH3)6 associated with d-tartaric acid produces a pronounced Cotton effect. These results show that some... 

Calculate the concentrations of molecular and ionic species in equilibrium with complex ions (Section 16.6, Problems 43-46). 

Figure 3.3 Fraction of Cd present as the free ion and as chloride complexes as a function of chloride ion concentration (M) for a total Cd concentration of 0.010 M at 25°C. Also shown are the distribution of species as computed in Example 3.2, and in seawater. Modified after J. N. Butler, 1964. Ionic equilibrium, a mathematical approach. Used by permission.

The choice of acid or base for solvent is simplified appreciably for melts containing complex ions (as a rule, they are anions), which are prone to the acid-base dissociation. Dissociation of this ion is assumed as the intrinsic acid-base equilibrium of a melt of such kind. In this case, the simpler eliminated anion will be considered as the base of the solvent and the coordinationally unsaturated residue will be the acidic particle of the solvent. Naturally, the division of particles formed by the auto-dissociation into acids and bases is made on the basis of the Lewis definition [13] an acid is the acceptor of an electron pair and a base is the donor of this electron pair. Ionic melts based on complex halides of gallium(III) [28], aluminium(III) [29] and boron(III) [30,31] may serve as examples of successful application of the above approach. The electron-deficient covalent halide (e.g. A1C13, BF3) in these melts is the solvent acid, and the corresponding halide ion is the base of the solvents ... 

Consider just a few cases of aqueous equilibria. The magnificent formations i n limestone caves and the vast expanses of oceanic coral reefs result from subtle shifts in carbonate solubility equilibria. Carbonates also influence soil pH and prevent acidification of lakes by acid rain. Equilibria involving carbon dioxide and phosphates help organisms maintain cellular pH within narrow limits. Equilibria involving clays in soils control the availability of ionic nutrients for plants. The principles of ionic equilibrium also govern how water is softened, how substances are purified by precipitation of unwanted ions, and even how the weak acids in wine and vinegar influence the delicate taste of a fine French sauce. In this chapter, we explore three aqueous ionic equilibrium systems acid-base buffers, slightly soluble salts, and complex ions. 

The final type of aqueous ionic equilibrium we consider involves a different kind of ion than we ve examined up to now. A simple ion, such as Na" " or S04 , consists of one or a few bound atoms, with an excess or deficit of electrons. A complex ion consists of a central metal ion covalently bonded to two or more anions or molecules, called ligands. Hydroxide, chloride, and cyanide ions are some ionic ligands water, carbon monoxide, and ammonia are some molecular ligands. In the complex ion Cr(NH3)6, for example, Cr is the central metal ion and six NH3 molecules are the ligands, giving an overall 3-1- charge (Figure 19.13). 

In several of the studies of aqueous chemistry of aluminum that have been made since about 1950, polynuclear complexing mechanisms have been proposed to identify and describe the dissolved aluminum hydroxide complex species (3, JO, 11). The formulae proposed have generally been based on stoichiometric considerations and pH measurements assuming the polynuclear species were ionic, and that equilibrium was attained. The complex ions reported by Hsu and Bates (8) were single six-mem-bered rings Ale (OH) 12 or multiples of this unit. Johansson (JO) identified a structural unit containing 13 aluminum and 40 oxygen atoms with various numbers of protons in crystalline basic aluminum sulfate. Because this solid formed readily, the same structural unit of aluminum was proposed as a solute species. Most of the proposed formulae for polynuclear complexes, however, have not been derived from structural considerations. 

In this experiment you will study qualitatively a complex ion equilibrium, and you will determine the equilibrium constant for the formation of FeNCS and see if the value stays the same as you change the concentration. The net ionic equation that describes the reaction is... 

The solubility product principle can only be strictly applied to equilibrium conditions, although it has often been used to explain such precipitations as those encountered in qualitative analysis by the traditional wet-test methods. However, these sudden precipitations do not take place under anything like equilibrium conditions and the fact that reasonably successful predictions can usually be made is mainly due to the enormous excess ionic concentrations (supersaturations) generated compared with those required by the corresponding solubility products. Errors of magnitude of 10 —10 per cent have been estimated (Lewin, 1960) for such calculations and these clearly swamp other variations such as neglect of solute activity coefficients, complex ion formation, etc. 

Complex composition of catalytic systems Many real catalytic systems are complex. An ionic polymerisation process may form the following equilibrium structures in catalytic systems two contact ionic pairs, solvent-separated ionic pairs, free ions, and aggregated particles. The presence of AC, of various reactivities, will influence the molecular characteristics of the forming polymers, especially their MWD curves. 

Two reports have appeared of kinetic studies of the base hydrolysis of complex ions of the type [Co(NH3)5(OCOR)] +. Activation parameters are reported from rate constants extrapolated to zero ionic strength for R = CM3 CI (/i = 1,2, or 3). When RCO2" = salicylate ion a simple second-order rate law is not observed owing to the importance of the deprotonation of the phenolic group. " If K is the equilibrium constant for this deprotonation reaction, the pseudo-first-order rate constant (k) in the presence of an excess of OH" ion is given by k=(A iis [OH-] k2K[OH-] )IH J5 [OH"])... 

Ballatori N, Shi C, Boyer JL (1988) Altered plasma membrane ion permeability in mercury-induced cell injury studies in hepatocytes of elasmobranch Raja erinacea. Toxicol Appl Pharmacol 95 279-291 Benndorf K, Nilius B (1988) Different blocking effects of Cd and Hg on the early outward current in myocardial mouse cells. Gen Physiol Biophys 7 345-352 Blazka ME, Shaikh ZA (1991) Differences in cadmium and mercury uptakes by hepatocytes role of calcium channels. Toxicol Appl Pharmacol 110 355-363 Brunder DG, Dettbarn C, Palade P (1988) heavy metal-induced Ca " release from sarcoplasmic reticulum. J Biol Chem 263 18785-18792 Butler JN (1964) Introduction to complex formation equilibria. Ionic equilibrium, a mathematical approach. Addison-Wesley, Reading, Massachusetts Palo Alto London, p 261... 

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