Redox potentials cobalt complexes

Similar effects are observed in the iron complexes of Eqs. (9.13) and (9.14). The charge on the negatively charged ligands dominates the redox potential, and we observe stabilization of the iron(iii) state. The complexes are high-spin in both the oxidation states. The importance of the low-spin configuration (as in our discussion of the cobalt complexes) is seen with the complex ions [Fe(CN)6] and [Fe(CN)6] (Fq. 9.15), both of which are low-spin. 

Jensen was the first to report in 1983 that the color of the solution oscillated between pink and dark brown in the presence of cobalt(II) and bromide ions when the reaction was carried out in a 90/10 (w/w) acetic acid/water mixture (162). This color change was accompanied by a change in the redox potential and the oscillations were observed for over 16 h and 800 cycles. Presumably, the pink color corresponds to a low Co(III)/Co(II) ratio, the dark brownish black to a high Co(III)/Co(II) ratio or to a Co(III)Br complex in this reaction. 

Table 1 lists some of the binding constants and rate constants measured for the reaction of CO2 with redox-active molecules. Various techniques have been used to measure these constants including cyclic voltammetry, pulsed radiolysis, and bulk electrolysis followed by UV-visible spectral measurements. The binding constants span an enormous range from less than 1 to 10 M [13-17]. Co(I) and Ni(I) macrocyclic complexes have been studied in some detail [13-16]. For the cobalt complexes, the CO2 binding constants K) and second-order rate constants for CO2 binding (kf) are largely determined by the Co(II/I) reduction potentials... 

We are still further from being able to explain the anodic activity of the CoTAA complex. The cobalt phthalocyanine, which is structurally identical with CoTAA in the inner coordination sphere, is completely inactive in the catalysis of anodic reactions. It therefore looks as if the central region is not exclusively responsible for the anodic activity. On the other hand, the fact that CoTAA is inactive for the oxidation of H2 points to n orbitals of the fuel participating in the formation of the chelate-fuel complex. A redox mechanism (cf. Section 5.2) can be ruled out because anodic oxidation proceeds only in the region below the redox potential of CoTAA (i.e. at about 600—650 mV). 

The resultant hydroxyl radicals are effective in initiating many chain reactions. The number of metal ions and complexes which are capable of activating hydrogen peroxide in this manner is quite large and is determined in part by the redox potentials of the activator. Related systems in which free radicals are generated by the intervention of suitable metallic catalysts include many in which oxygen is consumed in autoxidations. Cobalt(H) compounds which act as oxygen carriers can often activate radicals in such systems by reactions of the type ... 

Its oxidation potential is such that it is oxidized to the kinetically inert 2 1 cobalt(III) complex (equation 6). The relevant redox potentials have not been measured, but support is provided by the isolation by Wittwer46 of 2 1 cobalt(II) complexes of tridentate azo compounds. These are stable under only a very limited range of conditions and readily undergo oxidation to the corresponding 2 1 cobalt(III) complexes. 

The redox potentials and the strain energies at the cobalt(III) and cobalt(II) oxidation states of die most stable conformers of a number of hexaaminecobalt(III/II) complexes are listed in Table 10.1. The strain energy difference between the two oxidation states was found to correlate with the experimentally determined reduction potential11331. Fig. 10.2 shows a plot of the redox potentials of the hexaaminecobalt(III/II) complexes from Table 10.1 as a function of die strain energy differences between the oxidized and reduced forms. The experimentally determined redox potentials are given as solid points while the line corresponds to the calculated potentials. Based on Eq. 10.1,... 

Since these reactions are influenced by changes in the redox potential of the metal complex, it is possible to change from one process to the microscopic reverse process by changing the ligands attached to the metal. For example, with acetate ligands cobalt(II) is stable with respect to cobalt(III), and, in the presence of bromide ions, cobalt(III) is reduced by alkyl radicals in a ligand transfer oxidation ... 

The ESR spin Hamiltonian parameters and the redox potentials for the BDHC complexes are similar in magnitude to those for the corresponding corrinoid complexes. Thus, BDHC and corrinoid complexes are similar in the electronic nature of their nuclear cobalt. 

Table VI. Redox Potentials (V vs. SCE) for Various Cobalt Complexes...

The reaction of a Co(I) nucleophile with an appropriate alkyl donor is used most frequently for the formation of a Co-C bond, which also can be formed readily by addition of a Co(I) complex to an acetylenic compound or an electron-deficient olefin (5). The nu-cleophilicity of Co(I) in Co(I)(BDHC) is expected to be similar to that in the corrinoid complex, as indicated by their redox potentials. The formation of Co-C a-bond is the attractive criterion for vitamin Bi2 models. Sodium hydroborate (NaBH4) was used for the reduction of Co(III)(CN)2(BDHC) in tetrahydrofuran-water (1 1 or 2 1 v/v). The univalent cobalt complex thus obtained, Co(I)(BDHC), was converted readily to an organometallic derivative in which the axial position of cobalt was alkylated on treatment with an alkyl iodide or bromide. As expected for organo-cobalt derivatives, the resulting alkylated complexes were photolabile (17). 

The redox potentials and the strain energies at the cobalt(III) and cobalt(II) oxidation states of the most stable conformers of a number of hexaaminecobalt(III/ II) complexes are listed in Table 11.1 (selected data from[231]). The strain energy... 

The two cobalt(III) hexaamines with relatively short metal-donor distances considered here are [Co(7ra s-diammac)]3+ trans-diammac is trans-1,4,8,11 -tetraazacy-clotetradecane-6,13-diamine, see Fig. 17.11.1) and [Co(trap)2]3+ (trap is 1,2,3-tri-aminopropane, see Fig. 17.11.2). There are three conformers of [Co(taa s-di-ammac)]3+ (35, 51,15 where 5 and 1 refer to the conformation of the five-mem-bered chelate rings in the complex - see also Section 17.3). The A<5-conformer is the most stable form and has been characterized by an X-ray diffraction study. 15-[Co(fra .v-diammac)]3+ has very short Co-N bonds (1.937A (four equatorial bonds), 1.946 A (two axial bonds)), and the experimentally determined high ligand field and the strongly negative redox potential confirm that these structural features are conserved in solution190 231,2811. 

Redox potentials for a group of selected bis(dithiolene) complexes are listed in Table II. Neutral iron and cobalt bis(dithiolene) complexes exist in the dimeric form. The dimers stay intact when partially reduced and dissociate into monomers when fully reduced (34, 35). The potentials listed in Table II for Fe and Co complexes are therefore for the redox couples (0/—1, — 1/—2, etc.) of the dimer (Eq. 2). 

Redox Potential for Iron and Cobalt Nitrosyl Dithiolene Complexes, [M(NO)(S—S)2]z ... 

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