Copper complexes equilibrium constant

RAM/FAN] Ramette, R. W., Fan, G., Copper(II) chloride complex equilibrium constants, Inorg. Chem., 22, (1983), 3323-3326. Cited on page 413. 

There are a few documented examples of studies of ligand effects on hydrolysis reactions. Angelici et al." investigated the effect of a number of multidentate ligands on the copper(II) ion-catalysed hydrolysis of coordinated amino acid esters. The equilibrium constant for binding of the ester and the rate constant for the hydrolysis of the resulting complex both decrease in the presence of ligands. Similar conclusions have been reached by Hay and Morris, who studied the effect of ethylenediamine... 

Shuman and Michael [10] applied a rotating disk electrode to the measurement of copper complex dissociation rate constants in marine coastal waters. An operational definition for labile and non-labile metal complexes was established on kinetic criteria. Samples collected off the mid-Atlantic coast of USA showed varying degrees of copper chelation. It is suggested that the technique should be useful for metal toxicity studies because of its ability to measure both equilibrium concentrations and kinetic availability of soluble metal. 

Rate and equilibrium constant data, including substituent and isotope effects, for the reaction of [Pt(bpy)2]2+ with hydroxide, are all consistent with, and interpreted in terms of, reversible addition of the hydroxide to the coordinated 2,2 -bipyridyl (397). Equilibrium constants for addition of hydroxide to a series of platinum(II)-diimine cations [Pt(diimine)2]2+, the diimines being 2,2 -bipyridyl, 2,2 -bipyrazine, 3,3 -bipyridazine, and 2,2 -bipyrimidine, suggest that hydroxide adds at the 6 position of the coordinated ligand (398). Support for this covalent hydration mechanism for hydroxide attack at coordinated diimines comes from crystal structure determinations of binuclear mixed valence copper(I)/copper(II) complexes of 2-hydroxylated 1,10-phenanthroline and 2,2 -bipyridyl (399). 

The equilibrium between Mg2+ and the edta anion L4 can be compared to that between NH3 and H+, because the corresponding equilibrium constants are very dose in magnitude. In the latter case, it is possible to titrate NH3 with a solution of a strong acid in order to determine quantitatively its total concentration. It is therefore quite evident that, based on the values of the equilibrium constants of Scheme 3, the quantitative determination of the Mg2+ using edta should be possible. Because the other cations form more stable complexes than Mg2+, the complexometric titration should be of wide application. Some caution is necessary concerning the pH value at which the determination is done, because the ligand can be protonated, with consequent decrease of its chelating power. However, in the case of copper(II), its edta complex is already completely formed at pH 3 and therefore a titration is possible under these conditions. 

The equilibrium constant of reaction (1), K = [Cu ][Cu ]/[Cu ], is of the order of 10 thus, only vanishingly small concentrations of aquo-copper(I) species can exist at equilibrium. However, in the absence of catalysts for the disproportionation—such as glass surfaces, mercury, red copper(I) oxide (7), or alkali (311)—equilibrium is only slowly attained. Metastable solutions of aquocopper(I) complexes may be generated by reducing copper(II) salts with europium(II) (113), chromium(II), vanadium(II) (113, 274), or tin(II) chloride in acid solution (264). The employment of chromium(II) as reducing agent is best (113), since in most other cases further reduction to copper metal is competitive with the initial reduction (274). 

For each atom of copper deposited, three cyanide ions are released at the electrode surface. The concentration of free (CN) ions at the electrode surface is thus higher than its bulk concentration. This shifts the potential on the solution side of the double layer in the negative direction, lowering the concentration of the complex ions, hence lowering the rale of reaction. A typical copper cyanide bath is composed of 0.3 M CuCNandO.7 M KCN. The equilibrium constant in the reaction... 

The difference between die two slurries in polish rate behavior with slurry flow is a result of the degree of copper complexing obtained in each slurry. The ratio of ion concentration to Cu(NH3)2 ion concentration at equilibrium may be determined from the equilibrium constant for reaction (7.2) ... 

Complexes with Acetate and Other Common Brensted Bases Many reactions of copper tend to be conducted in the range of pH 4.0-5.5 to avoid possible formation of hydroxycopper species. This is a range in which acetate is the most commonly used buffer. Unfortunately, acetate (Ac ) is a very poor choice as it forms stronger complexes with Cu(II) than with any other divalent metal ion except Hg(II). The stepwise equilibrium constants for the formation of Cu(Ac)+, Cu(Ac)2, Cu(Ac)3 , and Cu(Ac)/ are 50, 10, 2.5, and 0.6, respectively ]176]. 

XAES) shows that CuLj chelates (LH = glutamine or asparagine) have distorted octahedral structures. The oxidative decarboxylation product of Gly-Gly-L-His and Cu(OH)2 is four-co-ordinate and square planar decarboxylation occurs at C-5 with deprotonation at C-4 to give a C==C system. Amine adducts of bis(ben-zoyl-/5-alaninato)copper(ii) have been isolated and characterized. Equilibrium constants have been reported for copper(ii) complexes of histidylhistidine from pH-titration data. " Qualitative analysis of the Cotton effect of d-d transitions of copper(ii) complexes with optically active acids has been obtained. ... 

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 ... 

---