Kinetics and mechanisms of the reactions of mononuclear and binuclear ruthenium(II) ammine complexes with peroxydisulfate

Abstract

Kinetic studies of the reactions of Ru(NH3)5pz2+ (pz = pyrazine), Ru(NH3)4bpy2+ (bpy = 2,2'-bipyridine), Ru(NH3)5pzMe3+ (pzMe+ = JV-methylpyrazinium), Ru2(NH3)10pz5+, RuRh(NH3)10pz5+, and Ru2(NH3)10pz4+ with S2Og1 2" are interpreted on the basis of a mechanism involving ion-pair formation between the ruthenium complexes and S2Os2" followed by one-electron transfer from Ru(II) to S2Os2". The dependence of the rate constants upon the standard free energy changes for the above reactions as

Discrete Cyanide‐Bridged Mixed‐Valence Co/Fe Complexes: Outer‐Sphere Redox Behaviour

“…10 Ϫ18  Ϫ1 s Ϫ1 by the direct application of the Marcus relationship. [23] Furthermore, no dependence on the pH is expected for the reduction potential in the acidity region used for this study. Given that some redox kinetics on dinuclear Ru/Fe mixed-valence compounds had already been studied with this oxidising agent, [24] Figure 2 shows the general kinetic behaviour found for these systems at varying [S 2 O 8 2Ϫ ], temperature and pH; second-order rate constants extrapolated at 298 K, and thermal and pressure activation parameters are collected in Table 3.…”

Salt effects upon the S2O82? + Ru(NH3)5pz2+ electron transfer reaction

“…Table II also contains the calculated redox potentials for the S 2 O 8 2Ϫ /S 2 O 8 3Ϫ couple. These values have been estimated from the values of the redox potential in water for this couple, 1.39 V [6], and using the measured values for IrCl 6 2Ϫ /IrCl 6 3Ϫ , assuming that the variations of redox potentials of the two anionic 2Ϫ/ 3Ϫ couples are the same in the salt solutions. The absence of significant specific salt effects on k obs and EЊЈ gives some support to this approximation.…”

“…The initial concentration of Ru(NH 3 ) 5 pz 2ϩ was 1.35 ϫ 10 Ϫ5 mol dm Ϫ3 and the S 2 O 8 2Ϫ concentration was 1.90 ϫ 10 Ϫ3 mol dm Ϫ3 . To avoid the protonation of the ruthenium complex [6] all the measurements were recorded at a fixed pH of 4.3 using an acetate/acetic acid buffer ([CH 3 COONa] ϭ 3.1 ϫ 10 Ϫ2 mol dm Ϫ3 and [CH 3 COOH] ϭ 7 ϫ 10 Ϫ2 mol dm Ϫ3 ). In all the experiments, temperature was maintained at 298.2 Ϯ 0.1 K. Pseudo-first-order rate constants were obtained from the slope of the plots of ln(A t Ϫ A ϱ ) vs time, where A t and A ϱ were the absorbances at time t and when the reaction was finished.…”

“…In this regard it can be shown that, in electron transfer reactions, this condition is accomplished by processes showing high internal free energy of reorganization, that is, by processes that are in the socalled slow reaction limit of Sumi and Marcus [9]. In that sense, peroxodisulphate oxidations are the ideal candidates, owing to the enormous internal free energy of reorganization of this oxidant [6]. In fact, this is the cause of the relatively slow reaction rates observed in peroxodisulphate oxidations, in spite of its high redox potential (see the following).…”

Kinetic studies on the metal(II) tartarate–peroxomonosulfate reaction

“…The chemistry of this oxidant has gained much attention in the recent past since the radical intermediates produced/derived is more reactive than the PMS itself [2]. In acidic and weakly alkaline pH, PMS with easily oxidizable metal ions such as Co(II) gives sulfate ion radical SO −• 4 ,which is a better oxidizing agent with a reduction potential of ∼ +2.4 to +2.6 V [3,4]. Some transition metal ion-PMS systems generate hydroxyl radical.…”