Vanadium complexes redox reactions

Vanadium(n). An interesting feature of many vanadium(ii) redox reactions is the similarity of kinetic data k, Aift, A t) to those for V + substitutions and the possibility that these processes are substitution controlled. The reaction with the oxalatotetramminecobalt(m) complex has been studied, the rate law being... 

A ubiquitous characteristic of vanadium chemistry is the fact that vanadium and many of its complexes readily enter into redox reactions. Adjustment of pH, concentration, and even temperature have often been employed in order to extend or maintain system integrity of a specific oxidation state. On the other hand, deliberate attempts to use redox properties, particularly in catalytic reactions, have been highly successful. Vanadium redox has also been successfully utilized in development of a redox battery. This battery employs the V(V)/V(IV) and V(III)AT(II) redox couples in 2.5 M sulfuric acid as the positive and negative half-cell electrolytes, respectively. Scheme 12.2 gives a representation of the battery. The vanadium components in both redox cells are prepared from vanadium pentoxide. There are two charge-discharge reactions occurring in the vanadium redox cells, as indicated in Equation 12.1 and Equation 12.2. The thermodynamics of the redox reactions involved have been extensively studied [8],... 

Also, coordination compounds and metal carbonyls are able to undergo a PET, resulting in initiating radicals [63]. Recently investigated examples are iron chloride based ammonium salts [149], vanadium(V) organo-metallic complexes [150], and metal sulfoxide complexes [151]. However, the polymerization efficiency of some systems is only low due to redox reactions between the central metal ion and the growing polymer radical, and the low quantum yields of PET. 

The redox reactivity of divanadium salen615 and other complexes555 have been investigated. Disproportionation of [ V(salen) 2(//-0)] was observed under electrochemical conditions.6 The reduced, [Vin(salen)]+ complex was found to be the essential species in the catalysis of the electroreduction of 02 by four electrons in CH2C12.616 While [ V(salcn) 2(/x-0)] was proposed to be the active species in the redox reaction, more recently [Vin(salen)]+ was identified as a reservoir from which the active species forms.616 Schiff base complexes encapsulated in zeolite Y are catalytically active in the oxidation of thioanisole with H202.14 Acid-promoted disproportionation of a Vlv phenolate under anaerobic conditions was proposed as a model reaction for the vanadium uptake in tunicates.139... 

Two tetrahedral complexes and one octahedral complex of vanadia species were found when vanadium and titanium were deposited simultaneously. Tetrahedral complexes are less active in redox reactions. All complexes differ from those found on the V2O5 on Si02 and the V2O5 on Ti02 catalysts. 

Two studies have been made of the oxidation of hypophosphorous acid. In perchlorate media the rate law of the redox reaction is indicative of several vanadium(v)-phosphorous(i) complexes being kinetically... 

A ternary vanadium(v)-edta-hydrogen peroxide complex of the type [V(edta)(Ha02)] has been characterized in solutions of pH 1.0—3.4 with little evidence for the occurrence of an intramolecular redox reaction. The reaction with neptunium(v) has been shown to result in an equilibrium between Np and HgOg slowly reduces Np " to Np in strongly... 

The action of redox metal promoters with MEKP appears to be highly specific. Cobalt salts appear to be a unique component of commercial redox systems, although vanadium appears to provide similar activity with MEKP. Cobalt activity can be supplemented by potassium and 2inc naphthenates in systems requiring low cured resin color lithium and lead naphthenates also act in a similar role. Quaternary ammonium salts (14) and tertiary amines accelerate the reaction rate of redox catalyst systems. The tertiary amines form beneficial complexes with the cobalt promoters, faciUtating the transition to the lower oxidation state. Copper naphthenate exerts a unique influence over cure rate in redox systems and is used widely to delay cure and reduce exotherm development during the cross-linking reaction. 

The experimental observations were interpreted by assuming that the redox cycle starts with the formation of a complex between the catalyst and the substrate. This species undergoes intramolecular two-electron transfer and produces vanadium(II) and the quinone form of adrenaline. The organic intermediate rearranges into leucoadrenochrome which is oxidized to the final product also in a two-electron redox step. The +2 oxidation state of vanadium is stabilized by complex formation with the substrate. Subsequent reactions include the autoxidation of the V(II) complex to the product as well as the formation of aVOV4+ intermediate which is reoxidized to V02+ by dioxygen. These reactions also produce H2O2. The model also takes into account the rapidly established equilibria between different vanadium-substrate complexes which react with 02 at different rates. The concentration and pH dependencies of the reaction rate provided evidence for the formation of a V(C-RH)3 complex in which the formal oxidation state of vanadium is +4. 

In this paper selectivity in partial oxidation reactions is related to the manner in which hydrocarbon intermediates (R) are bound to surface metal centers on oxides. When the bonding is through oxygen atoms (M-O-R) selective oxidation products are favored, and when the bonding is directly between metal and hydrocarbon (M-R), total oxidation is preferred. Results are presented for two redox systems ethane oxidation on supported vanadium oxide and propylene oxidation on supported molybdenum oxide. The catalysts and adsorbates are stuped by laser Raman spectroscopy, reaction kinetics, and temperature-programmed reaction. Thermochemical calculations confirm that the M-R intermediates are more stable than the M-O-R intermediates. The longer surface residence time of the M-R complexes, coupled to their lack of ready decomposition pathways, is responsible for their total oxidation. 

This book does not follow a chronological sequence but rather builds up in a hierarchy of complexity. Some basic principles of 51V NMR spectroscopy are discussed this is followed by a description of the self-condensation reactions of vanadate itself. The reactions with simple monodentate ligands are then described, and this proceeds to more complicated systems such as diols, -hydroxy acids, amino acids, peptides, and so on. Aspects of this sequence are later revisited but with interest now directed toward the influence of ligand electronic properties on coordination and reactivity. The influences of ligands, particularly those of hydrogen peroxide and hydroxyl amine, on heteroligand reactivity are compared and contrasted. There is a brief discussion of the vanadium-dependent haloperoxidases and model systems. There is also some discussion of vanadium in the environment and of some technological applications. Because vanadium pollution is inextricably linked to vanadium(V) chemistry, some discussion of vanadium as a pollutant is provided. This book provides only a very brief discussion of vanadium oxidation states other than V(V) and also does not discuss vanadium redox activity, except in a peripheral manner where required. It does, however, briefly cover the catalytic reactions of peroxovanadates and haloperoxidases model compounds. 

Hiatt et a/.34a-d studied the decomposition of solutions of tert-butyl hydroperoxide in chlorobenzene at 25°C in the presence of catalytic amounts of cobalt, iron, cerium, vanadium, and lead complexes. The time required for complete decomposition of the hydroperoxide varied from a few minutes for cobalt carboxylates to several days for lead naphthenate. The products consisted of approximately 86% tert-butyl alcohol, 12% di-fe/T-butyl peroxide, and 93% oxygen, and were independent of the catalysts. A radical-induced chain decomposition of the usual type,135 initiated by a redox decomposition of the hydroperoxide, was postulated to explain these results. When reactions were carried out in alkane solvents (RH), shorter kinetic chain lengths and lower yields of oxygen and di-te/T-butyl peroxide were observed due to competing hydrogen transfer of rm-butoxy radicals with the solvent. 

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