Redox cyanide complexes

The redox chemistries of di-, tri and tetra-nuclear Mo—S cyanide complexes have been discussed in relation to their electronic structures.135 Passage of oxygen into an aqueous solution of [Mo2S2(CN)s]6 leads to the formation of a dark violet mixed-crystal compound of the composition K4+J,[Mo2(S02)(S2)(CN)8] [Mo2(S02)(S2)(CN)8]1 -4H20 (x = 0.3). In the crystal the two anions, whilst structurally similar, are located at crystallographically independent positions each involves both molybdenum atoms surrounded by an approximately... 

The existence of Tl cyanide complexes has been mentioned previously however, on the basis of the analogy with other Tl -pseudohalide redox reactions, and by analysis of the redox potentials, the existence of the T1(CN) "" complexes was not widely accepted for a time. However, a detailed investigation of this system using ° T1 and NMR spectroscopy has indicated that Tl indeed forms very stable cyanide complexes (the overall formation constants,... 

An impressive number of articles on the redox kinetics of octacyano complexes have been produced during the past two decades. The material in this chapter covers the period between 1969 and 1991. Interested readers may find a good deal on the relatively few older mechanistic studies in reviews on mechanisms of redox reactions (37) and cyanide complexes of the early transition metals (7). A book by Sharpe (2) on... 

The use of ozone os on oxidant for industrial wastes containing cyanides and other reducible toxic substances appears worthy of careful investigation. The oxidation of simple cyanides by ozone is rapid and complete. Mass transfer controls the absorption. The use of packed towers or sieve plate towers is indicated, and the maintenance of a pH of at least 9.0 is recommended. The destruction of cyanates and cyanide complexes is slower than the cyanide oxidation. These substances are destroyed if sufficient contact time and proper pH control are maintained so that these slower reactions can take place. The use of redox potential to control the degree of oxidation appears promising. Proper interpretation of the redox potential of the treated waste will give an excellent indication of the effectiveness of the treatment and the degree of removal of cyanide and cyanate. 

Introduction.—The organometallic chemistry of technetium and rhenium reported during 1974 has been surveyed. The ligand-induced redox reactions of rhenium halides have been reviewed, and the chemistry of cyanide complexes of Group Vila metals has received attention. The electrochemistry of technetium and rhenium has been the subject of two reviews, and a general monograph on the production, uses, and disposal of technetium has appeared. Recently published crystal structures of complexes of technetium and rhenium have been collated. ... 

The driving force for an ET reaction at the ITIES consists of two components, i.e., the difference of standard potentials of redox mediators and the interfacial potential drop [Eq. (14)]. The dependence of kf on AE° was studied for reactions between ZnPor+ in benzene and the series of similar cyanide complexes in water (26) ... 

Hydrometallurgy makes use of redox and complexation reactions for the recovery of valuable metals from ores. This is exemplified by the dissolution of gold metal in ores through oxidation by air and complexation by cyanide. 

Table 8 Data on the energies and redox potentials of some cyanide complexes, rate constants for their quenching of Rufbpyh in aqueous solution at ambient temperature.

Ruthenium hexacyanides have been used in these photoelectron transfer reactions because of the availability of the (Ru(II)/Ru(III)) redox system, which differs by one electron. For the cyanide complex, (RuCCNji ) " =0.86 V. An example... 

The redox reactions of cobalt cyanide complexes with borohydide ion have been reported,the reactive species in the reduction of [Co(CN)6] being the hydrolysed ion BHgOH". The overall reaction is rapid (A 10 1 mol s ) and is controlled by the formation of the hydrolysed species. 

The redox potentials of zinc-substituted phthalocyanines are shown to be linearly dependent on the total Hammett substituent constant.837 In 1987, Stillman and co-workers used the absorption and magnetic circular dichroism spectra of the zinc phthalocyanine and its 7r-cation-radical species to assign the observed bands on the basis of theoretical calculations. The neutral and oxidized zinc phthalocyanine complexes with cyanide, imidazole, and pyridine were used with the key factor in these studies the stability of the 7r-cation-radical species.838 The structure of zinc chloro(phthalocyaninato) has been determined and conductivity investigated.839... 

Iron(III) very readily forms complexes, which are commonly 6-coordinate and octahedral. The pale violet hexaaquo-ion [Fe(H20)6]3+ is only found as such in a few solid hydrated salts (or in their acidified solutions), for example Fe2(S04)3.9H20. Fe(C104)3.10H20. In many other salts, the anion may form a complex with the iron(III) and produce a consequent colour change, for example iron(III) chloride hydrate or solution, p. 394. Stable anionic complexes are formed with a number of ions, for example with ethanedioate (oxalate), C204, and cyanide. The redox potential of the ironll ironlll system is altered by complex formation with each of these ligands indeed, the hexacyanoferrate(III) ion, [Fe(CN)6]3. is most readily obtained by oxidation of the corresponding iron(II) complex, because... 

Table XVI shows a selection of stability constants and redox potentials for iron(II) and iron(III) complexes. This Table covers a wide range of the latter, showing how the relative stabilities of the iron(II) and iron(III) complexes are refiected in. B (Fe /Fe ) values. A more detailed illustration is provided by the complexes of a series of linear hexadentate hydroxypyridinonate and catecholate ligands, where again high stabilities for the respective iron(III) complexes are refiected in markedly negative redox potentials (213). The combination of the high stabilities of iron(III) complexes of hydrox5rpyridinones, as of hydroxamates, catecholates, and siderophores, and the low stabilities of their iron(II) analogues is also apparent in Fig. 8. Here redox potentials for hydroxypyranonate and hydroxypyridinonate complexes of iron are placed in the overall context of redox potentials for iron(III)/iron(II) couples. The -(Fe /Fe ) range for e.g., water, cyanide, edta, 2,2 -bipyridyl, and (substituted) 1,10-phenanthrolines is...

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