Electrode surfaces, decomposition

Germanium In situ STM studies on Ge electrodeposition on gold from an ionic liquid have quite recently been started at our institute [59, 60]. In these studies we used dry [BMIM][PF<3] as a solvent and dissolved Gel4 at estimated concentrations of 0.1-1 mmol 1 the substrate being Au(lll). This ionic liquid has, in its dry state, an electrochemical window of a little more than 4 V on gold, and the bulk deposition of Ge started several hundreds of mV positive from the solvent decomposition. Furthermore, distinct underpotential phenomena were observed. Some insight into the nanoscale processes at the electrode surface is given in Section 6.2.2.3. 

In principle, the oxidation of proceeds at an electrode potential that is more negative by about 0.7 V than the anodic decomposition paths in the above cases however, because of the adsorption shift, it is readily seen that practically there is no energetic advantage compared to CdX dissolution in competing for photogenerated holes. Similar effects are observed with Se and Te electrolytes. As a consequence of specific adsorption and the fact that the X /X couples involve a two-electron transfer, the overall redox process (adsorption/electron trans-fer/desorption) is also slow, which limits the degree of stabilization that can be attained in such systems. In addition, the type of interaction of the X ions with the electrode surface which produces the shifts in the decomposition potentials also favors anion substitution in the lattice and the concomitant degradation of the photoresponse. 

C/pB estimated by both electrical (Mott-Schottky) and optical (photocurrent voltammetry) methods in the media studied, for (11 l)-oriented ZnSe electrode surfaces. A different variation was observed for the (110) orientation at pH >6. At pH 0, for both (110) or (11 l)-oriented electrode surface, the flat band potential value was -1.65 V (SHE) and the measured potential stability range (no detected current) was -0.35 to +2.65 V (SHE). A comparison of band levels with the other II-VI compounds as well as decomposition levels of ZnSe is given in Fig. 5.6. 

Theoretically, at a low Uappl the counteraction would be expected to result in full polarization of the electrodes, i.e., would become equal to Eappl, so no current will be passed however, the actual pc,2 at the electrode surface is continuously diminished by diffusion of the Cl2 gas into the solution and so there results a residual current, i = (2 appl - E fR. The amount of the latter increases more or less gradually with increasing Uappl, because the actual pC 2 increases until it finally becomes 1 atm, where Cl2 gas starts to escape from the solution. In the meantime, the anode has been completely covered with Zn metal, so that [Zn] has become unity. In fact, E has now attained a constant maximum value, the so-called decomposition potential, where electrolysis really breaks through. Any further increase in app) would, according to first expectations, cause a linear current increase, i = ( app, - Edecomp )IR. However, Fig. 3.2 shows that the experimental current curve deviates more and... 

When the potential applied to a polarographic cell exceeds the decomposition potential of an electroactive species, its concentration at the surface of the mercury drop is immediately diminished. A concentration gradient is thereby established and more of that species diffuses from the bulk solution to the electrode surface (Fick s law of diffusion). The resulting current flow is proportional to the rate of diffusion which in turn is determined by the concentration gradient, i.e. 

This electrochemical decomposition requires about 1 V at the electrode surface. To drive the protons into the WO3 film, a proton-conducting electrolyte, typically... 

In addition to the universal concern for catalytic selectivity, the following reasons could be advanced to argue why an electrochemical scheme would be preferred over a thermal approach (i) There are experimental parameters (pH, solvent, electrolyte, potential) unique only to the electrode-solution interface which can be manipulated to dictate a certain reaction pathway, (ii) The presence of solvent and supporting electrolyte may sufficiently passivate the electrode surface to minimize catalytic fragmentation of starting materials. (iii) Catalyst poisons due to reagent decomposition may form less readily at ambient temperatures, (iv) The chemical behavior of surface intermediates formed in electrolytic solutions can be closely modelled after analogous well-characterized molecular or cluster complexes (1-8). (v)... 

The second approach is an adaptation of the voltammetry technique to the working environment of electrolytes in an operational electrochemical device. Therefore, neat electrolyte solutions are used and the working electrodes are made of active electrode materials that would be used in an actual electrochemical device. The stability limits thus determined should more reliably describe the actual electrochemical behavior of the investigated electrolytes in real life operations, because the possible extension or contraction of the stability window, due to either various passivation processes of the electrode surface by electrolyte components or electrochemical decomposition of these components catalyzed by the electrode surfaces, would have been... 

In addition to the above thermodynamic consideration, kinetics also play an important role in determining the anodic stability of these salts. For example, some salts whose decomposition products are polymeric moieties were found to passivate the electrode surface effectively." Therefore, although the intrinsic oxidation potentials for these anions were not as high ( 4.0 V), they showed stability up to 4.50 V in subsequent scans. It should be cautioned here, though, as the passivation was only observed on an inert electrode surface, whether similar passivations would occur on an actual cathode surface... 

In thf the complexes Ni2 ( u.- j -PhC2R)Cp2 (including R = Ph, C CPh) undergo irreversible oxidation processes near -I-0.7 V (vs SCE, FcH /FcH " -l-O.ll V, FcH/FcH" " -I-0.56 V) which results in the formation of deposits on the electrode surface. The anodic sweep indicates the presence of a reversible reduction near — 1.30 V attributed to a Ni2-centered reduction and the formation of [Ni2(M-7 -PhC2R)Cp2] Further reduction results in decomposition of the complexes, and the liberation of the alkyne or diyne ligand, as evidenced by two characteristic alkyne/diyne reductions at very negative potentials. ... 

Flow coulometry experiments were performed to study the reduction of U02 in nitric, perchloric, and sulfuric acid solutions [56]. The results of these studies show a single two-electron reduction wave attributed to the U02 /U + couple. The direct two-electron process is observed without evidence for the intermediate U02" " species because of the relatively long residence time of the uranium ion solution at the electrode surface in comparison to the residence time typically experienced at a dropping mercury working electrode. The implication here is that as the UO2 is produced at the electrode surface, it is immediately reduced to the ion. As the authors note a simplified equation for this process can be written, Eq. (7), but the process is more complicated. Once the U02" species is produced it experiences homogeneous reactions comprising Eqns (8) and (9) or (8) and (10) followed by chemical decomposition of UOOH+ or UO + to [49]. 

Three anodic partial reactions are considered active dissolution of two metals M and M with different kinetics in the absence of their ions in bulk solution and decomposition of water with the evolution of oxygen. The kinetics of the latter process is so slow on most corroding metals that only at very negative potentials can oxygen present in the solution be electroreduced and this eventually becomes limited by mass transport due to the limited solubility of oxygen in water. At even more negative potentials, hydrogen evolution takes place on the electrode surface. The cathodic reduction of some metal ions present on the electrode surface as a consequence of corrosion is also considered in Fig. 13(b). 

Cyclic voltammetry can (i) determine the electrochemical reversibility of the primary oxidation (or reduction) step (ii) allow the formal potential, E°, of the reversible process to be estimated (iii) provide information on the number of electrons, n, involved in the primary process and (iv) allow the rate constant for the decomposition of the M"+ species to be measured. Additional information can often be obtained if intermediates or products derived from M"+ are themselves electroactive, since peaks associated with their formation may be apparent in the cyclic voltam-mogram. The idealized behaviour illustrated by Scheme 1 is a relatively simple process more complicated processes such as those which involve further electron transfer following the chemical step, pre-equilibria, adsorption of reactants or products on the electrode surface, or the attack of an electrogenerated product on the starting material, are also amenable to analysis. 

Equations 20.54 and 20.55 result from the fact that the ongoing chemical reaction occurs at the electrode surface also since this does take place at the electrode, however, the decomposition product of B is immediately converted to D. This flux must be accounted for in the determination of the dimensionless current parameter for that iteration,... 

Due to the chemical potential difference for species in the electrolyte and the photoelectrode, and by virtue of the fact that the electrode can be run in forward and reverse bias configurations, a number of important processes at the interface can be discerned. In each case, we will be concerned with the energy required for the process under consideration to occur and its resulting effects on photoelectrode performance. We can think of these processes as being of four basic types chemisorption, the desired electron or hole charge transfer, surface decomposition and electrochemical ion injection. In the rest of the paper we will briefly summarize our present understanding of each. 

In the further scans, this first wave is not observed because in the aforementioned potential region, the electrode surface is no longer a pure platinum one but is a rearranged platinum hydroxide surface39. The results described in section 6.2 showed that the limiting-current plateau of the second oxidation wave (first scan) is controlled by transport of dithionite. This indicates that electron transfer from dithionite to PtOH and/or PtO is a much faster process than transport of dithionite towards the electrode. This is confirmed by the fact that in the further scans an identical limiting-current is obtained. The third oxidation wave in the first scan (second wave for the other scans) is attributed to the oxidation of sulphite described earlier. It is formed as a reaction product of the sodium dithionite oxidation and also of the homogeneous decomposition of sodium dithionite. Also in this case, a hysteresis effect is observed for the sulphite forward/back-ward sweep oxidation wave. 

In systems where ECL arises upon application of a single potential and reaction of a coreactant, as outlined in Eqs. (5) through (9), concentration distance profiles differ and depend on many more factors. In cases where a positive potential is applied and both the ECL chromophore and coreactant are oxidized at the electrode surface, concentrations of the two initial radical ion species will decay with increasing distance from the electrode, as will the concentration of the strong reducing agent formed upon decomposition of the coreactant. The zone where luminescence arises depends on relative rate constants for... 

The cyclic voltammetric experiment can give a great deal of information about the redox activity of a compound and the stability and accessibility of its reduced or oxidised forms. For a fully chemically reversible process, ipa must equal rpc, i.e. all of the material oxidised at the electrode surface on the forward scan must be re-reduced on the reverse scan (or vice versa). If this condition does not hold true, then the process may be partially reversible (rpc < ipa) or irreversible (rpc = 0). Observation of processes that are not fully reversible implies decomposition or chemical reaction of the reduced or oxidised species and the ratio of ipa to /p(. will show a strong dependence on scan rate since the reverse current is related to the lifetime of the redox-generated material. Note that processes that are chemically reversible (in the sense that the reduced and oxidised species are both stable) may not be electrochemically reversible (a term that relates to the relative rates of forward and back electron transfer). Electrochemically reversible processes are characterised by a separation between the forward and reverse potential peaks of exactly 59 mV. 

The potential window can be limited by the decomposition potential of a solute, not just a solvent. In particular, reactions of anodic oxidation of halides (Cl-, Br, and I-) on diamond are highly irreversible and have much higher overvoltage (for Cl, by 1 V) than on platinum or graphite electrodes [97, 123, 124], In all probability this is due to poor adsorption of intermediates, that is, Cl, Br, and I atoms, on the diamond electrode surface. We recall that the outer-sphere reactions discussed in Section 6.1 generally do not involve adsorption of intermediates and thus are not... 

It is well known that ACN reacts with active metals (Li, Ca) to form polymers [48], These polymers are products of condensation reactions in which ACIST radical anions are formed by the electron transfer from the active metal and attack, nucleophilically, more solvent molecules. Species such as CH3C=N(CH3)C=N are probably intermediates in this polymerization. ACN does not react on noble metal electrodes in the same way as with active metals. For instance, well-re-solved Li UPD peaks characterize the voltammograms of noble metal electrodes in ACN/Li salt solutions. This reflects a stability of the Li ad-layers that are formed at potentials above Li deposition potentials. Hence, the cathodic limit of noble metal electrodes in ACN solutions is the cation reduction process (either TAA or active metal cations). As discussed in the previous sections, with TAA-based solutions it is possible that the electrode surfaces remain bare. When the cations are metallic (e.g., Li+), it is expected that the electrode surfaces become covered with surface films originating from atmospheric contaminants reduction if the electrodes are polarized below 1.5 V (Li/Li+). As Mann found [13], in the presence of Na salts the polarization of metal electrodes in ACN solutions to sodium deposition potentials leads to solvent decomposition, with evolution of H2, CH4 and sodium cyanide (due to reaction with metallic sodium). 

These requirements hold for the films at both the positive and negative electrode surfaces. Thus, these surface films frequently comprise quite complex mixtures of reaction products and their presence affects the kinetic properties of charge transfer across the interface. It is the deviation of surface film s properties from meeting this set of ideal requirements that is the single most important cause of cell failure in a large fraction of cases. When the decomposition reactions occur, a small amount of active material must also be irreversibly consumed. 

When a typical nonactive material is employed, the anode only acts as an electron sink. In this particular case, the scheme representing the oxidation processes can be explained as shown in Fig. 4.2. The first process that needs to be considered is the mass transfer of the compounds from the bulk zone to the anodic one. The organic compounds can undergo direct oxidation on the electrode surface. This process can either be one-stage or multistage, and proceeds until the final oxidation product is generated (usually, carbon dioxide). At the same time, the decomposition of water... 

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