Solubility equilibria formation constants

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. 

The computation of formation constants is considered to be the most important aspect of equilibrium theory, since this knowledge permits a full specification of the complexation phenomena. Once this information is in hand, the formulator can literally define the system at a given temperature through the manipulation of solution-phase parameters to obtain the required drug solubility. 

The equilibrium has been studied in (presumedly) noninteracting medium both by spectrophotometric (40) and by solubility (41) techniques. In chlorobenzene (41), the relative formation constants for the octahedral forms of the M(II) chlorides of Mn, Co, Ni, Cu, and Zn are as shown in Table VII. The greater tendency of Ni(II) toward the octahedral configuration is shown by a factor of about 7 x 103 in the constant over that for Co(II). 

Problems in this chapter include some brainbusters designed to bring together your knowledge of electrochemistry, chemical equilibrium, solubility, complex formation, and acid-base chemistry. They require you to find the equilibrium constant for a reaction that occurs in only one half-cell. The reaction of interest is not the net cell reaction and is not a redox reaction. Here is a good approach ... 

If CL, CM and pH are kept the same for a series of different cations, the position of equilibrium in equation (6a) would depend on the value of the formation constant, (3 , for the metal complex in question and the extent of precipitation on its solubility product, K. The more stable the complex and the lower its solubility in water, the greater the extent of precipitation. Essentially there is a competition between cations M"+, and protons, H+, for the free ligand anion Ox. ... 

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. 

In order to verify the differences observed in 4f-5f complex stability, we used an additional solvent extraction method based on the competition between a soluble organic chelatant (TTA] and the aqueous soluble azide ions for binding the metal ions. The extraction equilibrium is shown by equation (2] and the distribution coefficients CD] of the metal at constant pH are correlated with the formation constants by equation (3]. 

Conventionally, equilibrium constants involving a solid compound are denoted as solubility constants rather than as formation constants of the solid. An index s to the equilibrium constant indicates that the constant refers to a solubility process, as shown in Eqs.(II.17) to (11.19)... 

The study of the variation of the solubility with the selenite concentration is stated to have been carried at the constant ionic strengths 0.01 and 0.3 M. How this was accomplished was not clear from the information in the paper. The data in the tables rather seem to indicate that the ionic strength varied and reached 0.03 and 0.5 M, respectively, in the solution with the highest selenite concentrations. The analysis of the data was made with an equilibrium model that comprised the solubility equilibrium and the formation of the complex 0(8003)2 T us the formation of CoSe03(aq) was not included in the model. The analysis led to values of the solubility product at the two ionic strengths that appear to be inconsistent with the value obtained from the solubility in water. This result together with the improbable model made the review reject the outcome of the equilibrium analysis. 

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. 

Silver ion forms a stepwise 1 2 complex with ethylenediamine (en) with formation constants of = 5.0 X 10 and Kf2= 1.4 X 10. Calculate the solubility of silver chloride in 0.100 M ethylenediamine. Also calculate the equilibrium concentrations of Aglen)" and Ag(en)2 . ... 

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