Redox potential of metal atoms

The observation of the atom reactivity is mostly restricted to metal atoms formed by reduction of monovalent ions, because they may be generated in one step in a single pulse regime. 

Similar observations on the oxidation of the thallium atom or on the reduction of T1+ have been made by pulse radiolysis. They are in agreement, as for silver, with the value determined from the electrode potential and the sublimation energy of the bulk metal into atoms, i.e. °(T1 /T1 ) = —1.9 Vnhe-Silver ions complexed by cyanide, ammonia, or EDTA, Ag L, are not reduced by the radical (CH3)2C OH, even under basic conditions, and the redox potential of these complexed forms must be more negative than —2.1 According 

Ion complexation by a strong ligand therefore induees a marked shift of the atom redox potential to more negative values compared with the aqueous system (Fig. 6), as is observed for the potentials of bulk metal electrodes. Similarly, the higher 

Because the sublimation energy of metals is usually large, the redox potential of all metal atoms ii°(M+/M°) is systematically quite negative (Table 4).i6- 7 3i.58] p j. instance, for copper °(Cu /Cu ) = E (Cu+/Cumet) - AGsub = -2.7 and 

Note that the quite negative redox potential of metal atoms °(M /M ) explains why the very first atom formation process from free or complexed monovalent ions in the bulk is thermodynamically unfavorable, unless the reducing agent is very strong. When the ions are adsorbed on a support, however, their reduction potential is markedly shifted to a higher value and reduction by a moderate electron donor is possible. For that reason, in the latter circumstance, the walls and any particle present in the solution play an important role in the nucleation. For example, the 

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In conclusion, the order of reduction of metal ions is controlled by their redox potential. This is also true in other pairs of precious metals such as Pd/Pt, Au/Pd, etc. (53). In addition, poly(jV-vinyl-2-pyrrolidone) (PVP) plays an important role for the formation of the core/shell structure. In the case of the Au/Pt system, the aggregation starts from Au but not Pt. This is probably due to the coordinating ability of metals to PVP. The Pt atoms or microclusters coordinating to PVP are more stable than the Au atoms or microclusters, since Au cannot coordinate to PVP. Thus, Au atoms or microcluster aggregate at first after the reduction, and then Pt atoms or microclusters deposit on the Au nuclei. In summary, the core/shell structure is controlled by (1) the redox potential of metal ions, and (2) the coordination ability of metals to PVP, stabilizing polymer. 

Our first attempt of a successive reduction method was utilized to PVP-protected Au/Pd bimetallic nanoparticles [125]. An alcohol reduction of Pd ions in the presence of Au nanoparticles did not provide the bimetallic nanoparticles but the mixtures of distinct Au and Pd monometallic nanoparticles, while an alcohol reduction of Au ions in the presence of Pd nanoparticles can provide AuPd bimetallic nanoparticles. Unexpectedly, these bimetallic nanoparticles did not have a core/shell structure, which was obtained from a simultaneous reduction of the corresponding two metal ions. This difference in the structure may be derived from the redox potentials of Pd and Au ions. When Au ions are added in the solution of enough small Pd nanoparticles, some Pd atoms on the particles reduce the Au ions to Au atoms. The oxidized Pd ions are then reduced again by an alcohol to deposit on the particles. This process may form with the particles a cluster-in-cluster structure, and does not produce Pd-core/ Au-shell bimetallic nanoparticles. On the other hand, the formation of PVP-protected Pd-core/Ni-shell bimetallic nanoparticles proceeded by a successive alcohol reduction [126]. 

The nature of the ligand donor atom and the stereochemistry at the metal ion can have a profound effect on the redox potential of redox-active metal ions. The standard redox potentials of Cu2+/Cu+, Fe3+/Fe2+, Mn3+/Mn2+, Co3+/Co2+, can be altered by more than 1.0 V by varying such parameters. A simple example of this effect is provided by the couple Cu2+/Cu+. These two forms of copper have quite different coordination geometries, and ligand environments, which are distorted towards the Cu(I) geometry, will raise the redox potential, as we will see later in the case of the electron transfer protein plastocyanin. 

Not all cytochromes from sulfate-reducing bacteria reduce Fe(III) or other metals. D. vulgaris produces a cyt C553, which has a molecular mass of 9 kDa, midpoint redox potential of OmV, and a single heme and the iron atom is coordinated by histidine methionine. It is unclear at this time if the inability of this cyt C553 to reduce metals is due to lack of a bishistidinyl iron coordination or to some other factor, such as steric hinderance owing to orientation of heme in the protein. 

It was also observed, in 1973, that the fast reduction of Cu ions by solvated electrons in liquid ammonia did not yield the metal and that, instead, molecular hydrogen was evolved [11]. These results were explained by assigning to the quasi-atomic state of the nascent metal, specific thermodynamical properties distinct from those of the bulk metal, which is stable under the same conditions. This concept implied that, as soon as formed, atoms and small clusters of a metal, even a noble metal, may exhibit much stronger reducing properties than the bulk metal, and may be spontaneously corroded by the solvent with simultaneous hydrogen evolution. It also implied that for a given metal the thermodynamics depended on the particle nuclearity (number of atoms reduced per particle), and it therefore provided a rationalized interpretation of other previous data [7,9,10]. Furthermore, experiments on the photoionization of silver atoms in solution demonstrated that their ionization potential was much lower than that of the bulk metal [12]. Moreover, it was shown that the redox potential of isolated silver atoms in water must... 

Moreover, almost in all the early steps, the redox potential of the clusters, which decreases with the nuclearity, is quite negative. Therefore the growth process undergoes another competition with a spontaneous corrosion by the solvent and the radiolytic protons, corrosion which may even prevent the formation of clusters, as mostly in the case of nonnoble metals. Monomeric atoms and oligomers of these elements are so fragile to reverse oxidation by the medium that H2 is evolved and the zerovalent metal is not formed [11]. For that reason, it is preferable in these systems to scavenge the protons by adding a base to the solution and to favor the coalescence by a reduction faster than the oxidation [53]. 

The value of the critical nuclearity allowing the transfer from the monitor depends on the redox potential of this selected donor S . The induction time and the donor decay rate both depend on the initial concentrations of metal atoms and of the donor [31,62]. The critical nuclearity corresponding to the potential threshold imposed by the donor and the transfer rate constant value, which is supposed to be independent of n, are derived from the fitting between the kinetics of the experimental donor decay rates under various conditions and numerical simulations through adjusted parameters (Fig. 5) [54]. By changing the reference potential in a series of redox monitors, the dependence of the silver cluster potential on the nuclearity was obtained (Fig. 6 and Table 5) [26,63]. 

The electrochemical behavior of tetracoordinated Cu(i) complexes (i.e., Cu(dpp)2-based cores) is well established.193,941 The reversible redox potential for the Cu(ii)/ Cu(i) transition is around 0.6-0.7 V versus SCE. This relatively high potential underlines the stability of the 4-coordinate Cu(i) complexes relative to their Cu(n) counterparts. The redox potential of pentacoordinated copper complexes 84 86 is observed in a much more cathodic range. For example, for the 5-coordinate complex Cu(l, dap)2+/+ (dap = 2,9-di-p-anisyl-l, 10-phcnanthrolinc), in which the terpy fragment of the ring is bound to the metal, the redox potential is -0.035 V. This potential shift when going from tetracoordinated to pentacoordinated copper systems is due to the better stabilization of the Cu(ii) state, thanks to the presence in the coordination sphere of live donor atoms. 

The redox potentials for the electron acceptors that react with HO (Table 15) are such that a pure outer-sphere single-electron transfer (SET) step would be endergonic (the HO /HO redox potential is more positive than the redox potential of the electron acceptor). Hence, the observed net reactions must be driven by coupled chemical reactions, particularly bond formation by the HO to the electrophilic atom of the acceptor molecule that accompanies a singleelectron shift. (The formation of the bond provides a driving force sufficient to make the overall reaction thermoneutral or exergonic 1.0 V per 23.1 kcalmol of bond energy.) The effect of various transition metal complexes on the oxidation potential for HO in MeCN illustrates some of these effects the results are summarized in Table 16. ... 

The standard reduction potentials (see Redox Potential) of the elements and their compounds have many important applied implications for chemists, not the least of which is being aware when a compound or mixture of compounds they are handling has the potential for exploding. This should be considered as a possibility when the appropriate potentials differ by more than about one volt and appropriate kinetics considerations apply. A simply predictable case is the sometimes-violent reaction of metals with acids, as illustrated in a recently produced discovery video. Redox activities of elements are most commonly (and most precisely) analyzed via thermo chemical cycles such as the familiar Bom-Haber cycle for the production of NaCl from Na and CI2. A similar analysis of the activities of different metals in their reactions with acids shows that the standard reduction potential for the metal (the quantitative measure of the activity of the metal) can be expressed in terms of the appropriate ionization energies of the metal, the atomization energies of the metal (see Atomization Enthalpy of Metals), and the hydration energies... 

As = surface area of a semiconductor contact [A ] = concentration of the reduced form of a redox couple in solution [A] = concentration of the oxidized form of a redox couple in solution A" = effective Richardson constant (A/A ) = electrochemical potential of a solution cb = energy of the conduction band edge Ep = Fermi level EF,m = Fermi level of a metal f,sc = Fermi level of a semiconductor SjA/A") = redox potential of a solution ° (A/A ) = formal redox potential of a solution = electric field max = maximum electric field at a semiconductor interface e = number of electrons transferred per molecule oxidized or reduced F = Faraday constant / = current /o = exchange current k = Boltzmann constant = intrinsic rate constant for electron transfer at a semiconductor/liquid interface k = forward electron transfer rate constant = reverse electron transfer rate constant = concentration of donor atoms in an n-type semiconductor NHE = normal hydrogen electrode n = electron concentration b = electron concentration in the bulk of a semiconductor ... 

Evidently, the redox potential values of clusters must be comprised between the limit values of the atom and the bulk metal. As the other physicochemical properties, redox potentials of clusters depend on their nuclearity, mostly at low n (Fig. Kinetic studies,... 

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