Chemical equilibrium complex ions

In Section 16-6, we describe how metal cations in aqueous solution can form bonds to anions or neutral molecules that have lone pairs of electrons. This leads to formation of complex ions and to chemical equilibria involving complexation. The complexation equilibrium between Ag and NH3 is an example ... 

Recently, Muha (83) has found that the concentration of cation radicals is a rather complex function of the half-wave potential the concentration goes through a maximum at a half-wave potential of about 0.7 V. The results were obtained for an amorphous silica-alumina catalyst where the steric problem would not be significant. To explain the observed dependence, the presence of dipositive ions and carbonium ions along with a distribution in the oxidizing strengths of the surface electrophilic sites must be taken into account. The interaction between the different species present is explained by assuming that a chemical equilibrium exists on the surface. 

The interaction between the ions is regarded as a chemical equilibrium. The resulting product of such an association between two individual kinds of ions is called a complex species, and the extent to which these various complex species occurs are described in terms of the equilibrium constants. 

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. 

Thus, for the conditions where second-order behavior is observed, the chemical circumstances indicate the cerium(IV) oxidation of each chromium complex will involve a rate-determining one-equivalent oxidation of the complex ion (or a species in rapid equilibrium with the complex ion) to an intermediate, followed by the rapid one-equivalent oxidation of the intermediate. Without reference to the role of water coordinated to the chromium, the most obvious mechanism in accord with these specifications is ... 

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. 

PK. A measurement of the complete ness of an incomplete chemical reaction. It is defined as the negative logarithm ito the base 101 of the equilibrium constant K for the reaction in question. The pA is most frequently used to express the extent of dissociation or the strength of weak acids, particularly fatty adds, amino adds, and also complex ions, or similar substances. The weaker an electrolyte, the larger its pA. Thus, at 25°C for sulfuric add (strong acid), pK is about -3,0 acetic acid (weak acid), pK = 4.76 bone acid (very weak acid), pA = 9.24. In a solution of a weak acid, if the concentration of undissociated acid is equal to the concentration of the anion of the acid, the pAr will be equal to the pH. 

In aquatic ecosystems, complexation to organic and inorganic ligands and competition between toxic metals and Ca or Mg ions for biological adsorption sites reduce the actual amount of metal available for uptake by organisms. Chemical equilibrium models applicable to natural systems include RANDOM (Murray and Linder 1983 ... 

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]. 

It is important to emphasize that the chemical behavior of the successor complex [ArH+, N02 ] in Class II systems is more like that of the pair of free radicals ArH+ and N02. The extent to which the N02 moiety in the successor complex is already bent (i.e. substantially) will further facilitate collapse to the cr-adduct. In this regard, the successor complex in Class II systems is very different from the Class III pre-equilibrium complex [ArH, NO]+, which exhibits ion-radical behavior only upon separation to the free ArH+ and NO species [61]. 

Dissociation — is the separation or splitting of a chemical compound (complexes, molecules, or salts) into two or more -> ions by dissolution and -> solvation, or by any other means, the breaking into smaller molecules, or radicals. In case of solvation, this results in an ioni-cally conducting -> electrolyte solution. D. usually occurs in a reversible manner. The opposite process is association or recombination. Assuming a reversible dissociation reaction in a chemical -> equilibrium of the form XY X + Y, the ratio of dissociation is quantified by the dissociation constant JCn, i.e JCn = where a denotes the activity of the species. The dissociation constants are frequently quoted as values of pAT = - log K. In mass spectrometry, the term is used in the meaning of a fragmentation, i.e., a decomposition of an ion into another ion of lower mass and one or more neutral species. 

Physico-chemical speciation refers to the various physical and chemical forms in which an element may exist in the system. In oceanic waters, it is difficult to determine chemical species directly. Whereas some individual species can be analysed, others can only be inferred from thermodynamic equilibrium models as exemplified by the speciation of carbonic acid in Figure 9. Often an element is fractionated into various forms that behave similarly under a given physical (e.g., filtration) or chemical (e.g., ion exchange) operation. The resulting partition of the element is highly dependent upon the procedure utilised, and so known as operationally defined. In the following discussion, speciation will be exemplified with respect to size distribution, complexation characteristics, redox behaviour and methylation reactions. 

Thermodynamic data, whether determined through calorimetry or solubility studies, are subject to refinement as more exact values for the components in the reaction scheme, or more complete description of the solution phases, become available. Many of the solubility studies on clays were done before digital-computer chemical equilibrium programs were available. One such program, SOLMNEQ, written by one of the authors ( ) solves the mass-action and mass-balance equations for over 200 species simultaneously. SOLMNEQ was employed in this investigation to convert the chemical analytical data into the activities of appropriate ions, ion pairs, and complexes. 

The use of the extended Debye-Hlickel equation with the appropriate equilibrium constants for mass action expressions to solve a complex chemical equilibrium problem is known as the ion-association (lA) method. 

Papelis C., Hayes K. F., and Leckie J. O. (1988) HYDRAQL A Program for the Computation of Chemical Equilibrium Composition of Aqueous Batch Systems Including Surface-complexation Modeling of Ion Adsorption at the Solution Oxide/Solution Interface. 306. Environmental Engineering and Science, Department of Civil Engineering, Stanford University, Stanford, CA, 131pp. 

Analytical Difficulties Although we can make many chemical equilibrium models that predict the existence of complexes in natural waters, analytically one encounters difficulties in identifying unequivocally the various solute species and in distinguishing between dissolved and particulate concentrations. The analytical task is rendered veiy difficult because the individual chemical species are often present at nano- and picomolar concentrations. The ion-selective electrode (ISE), if it were sufficiently sensitive, would permit the measurement of free metal-ion activity. 

In this mode of separation, active compounds that form ion pairs, metal complexes, inclusion complexes, or affinity complexes are added to the mobile phase to induce enantioselectivity to an achiral column. The addition of an active compound into the mobile phase contributes to a specific secondary chemical equilibrium with the target analyte. This affects the overall distribution of the analyte between the stationary and the mobile phases, affecting its retention and separation at the same time. The chiral mobile phase approach utilizes achiral stationary phases for the separation. Table 1 lists several common chiral additives and applications. 

All of the ions in the previous equation are colorless, so you cannot see which complex ion has the greater concentration. However, in the following chemical equilibrium of nickel ions, ammonia, and water, the complex ions have different colors. Therefore, you can see which complex ion has the greater concentration. 

The beaker on the left in Figure 10 contains this copper complex ion reaction in a chemical equilibrium that favors the formation of reactants. We know that the reverse reaction is favored, because the reaction mixture in the beaker is pale blue. But if additional ammonia is added to this beaker, the system responds to offset the increase by forming more of the product. This increase in the presence of product can be seen in the beaker to the right in Figure 10, which contains a blue-purple solution. 

According to IUPAC the definition solvents polarity is the overall solvation capability (or solvation power) for (1) educts and products, which influences chemical equilibrium, (2) reactants and activated complexes ( transition states ), which determines reaction rates, and (3) ions or molecules in their ground and first excited state, which is responsible for light absorptions in the various wavelength regions. This overall solvation capability depends on the action of all, non-specific and specific, intermolecular solute-solvent interactions, excluding such interactions leading to definite chemical alterations of the ions or molecules of the solute [53],... 

A review of some leading semiempirical models precedes examination of physicochemical modeling of ion exchange. Such models will likely be used for the foreseeable future to describe ion-exchange phenomena in complex svstems. Thus, they represent the reference point for development of improved models. Methods of incorporating the semiempirical ion-exchange equations in general chemical equilibrium models are also described. 

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