Standard potentials, electrolyte solutions

Standard potentials Ee are evaluated with full regard to activity effects and with all ions present in simple form they are really limiting or ideal values and are rarely observed in a potentiometric measurement. In practice, the solutions may be quite concentrated and frequently contain other electrolytes under these conditions the activities of the pertinent species are much smaller than the concentrations, and consequently the use of the latter may lead to unreliable conclusions. Also, the actual active species present (see example below) may differ from those to which the ideal standard potentials apply. For these reasons formal potentials have been proposed to supplement standard potentials. The formal potential is the potential observed experimentally in a solution containing one mole each of the oxidised and reduced substances together with other specified substances at specified concentrations. It is found that formal potentials vary appreciably, for example, with the nature and concentration of the acid that is present. The formal potential incorporates in one value the effects resulting from variation of activity coefficients with ionic strength, acid-base dissociation, complexation, liquid-junction potentials, etc., and thus has a real practical value. Formal potentials do not have the theoretical significance of standard potentials, but they are observed values in actual potentiometric measurements. In dilute solutions they usually obey the Nernst equation fairly closely in the form ... 

Several significant electrode potentials of interest in aqueous batteries are listed in Table 2 these include the oxidation of carbon, and oxygen evolution/reduction reactions in acid and alkaline electrolytes. For example, for the oxidation of carbon in alkaline electrolyte, E° at 25 °C is -0.780 V vs. SHE or -0.682 V (vs. Hg/HgO reference electrode) in 0.1 molL IC0 2 at pH [14]. Based on the standard potentials for carbon in aqueous electrolytes, it is thermodynamically stable in water and other aqueous solutions at a pH less than about 13, provided no oxidizing agents are present. 

The electrical double layer at pc-Zn/fyO interfaces has been studied in many works,154 190 613-629 but the situation is somewhat ambiguous and complex. The polycrystalline Zn electrode was found to be ideally polarizable for sufficiently wide negative polarizations.622"627 With pc-Zn/H20, the value of Eg was found at -1.15 V (SCE)615 628 (Table 14). The values of nun are in reasonable agreement with the data of Caswell et al.623,624 Practically the same value of Eff was obtained by the scrape method in NaC104 + HjO solution (pH = 7.0).190 Later it was shown154,259,625,628 that the determination of Eo=0 by direct observation of Emin on C,E curves in dilute surface-inactive electrolyte solutions is not possible in the case of Zn because Zn belongs to the group of metals for which E -o is close to the reversible standard potential in aqueous solution. 

The copper product is known as blister copper because of the appearance of air bubbles in the solidified metal. In the hydrometallurgical process, soluble Cu2+ ions are formed by the action of sulfuric acid on the ores. Then the metal is obtained by reducing these ions in aqueous solution either electrolytically or chemically, by using an inexpensive reducing agent that has a more negative standard potential than that of copper, such as hydrogen or iron (see Section 14.3) ... 

A standard rotating disk electrode (RDE) setup with a gas-tight Pyrex cell was used for the experiment on CO adsorption and the HOR. A Pt wire was used as counterelectrode. A reversible hydrogen electrode, RHE(t), kept at the same temperature as that of the cell (t, in °C), was used as the reference. All the electrode potentials in this chapter will be referenced to RHE(f). The electrolyte solution of 0.1 M HCIO4... 

Le Hung presented a general theoretical approach for calculating the Galvani potential Ajyj at the interface of two immiscible electrolyte solutions, e.g., aqueous (w) and organic solvent (s) [25]. Le Hung s approach allows the calculation of when the initial concentration (Cj), activity coefficients (j/,) and standard energies of transfer of ions (AjG ) are known in both solutions. 

As mentioned above, the distribution of the various species in the two adjacent phases changes during a potential sweep which induces the transfer of an ion I across the interface when the potential approaches its standard transfer potential. This flux of charges across the interface leads to a measurable current which is recorded as a function of the applied potential. Such curves are called voltammograms and a typical example for the transfer of pilocarpine [229] is shown in Fig. 6, illustrating that cyclic voltammograms produced by reversible ion transfer reactions are similar to those obtained for electron transfer reactions at a metal-electrolyte solution interface. 

In the chronopotentiogram of mixtures (see Fig. 3.58), the reactive components will yield different inflection points if their standard potentials show a sufficient difference (at least 0.1 V). To take a simple example, let us consider a reversible electrode process for both ox and ox2 in a solution of the supporting electrolyte. Then eqn. 3.72 is simply valid for the first reacting oxj with up to the first inflection point however, beyond this point the last traces of exj... 

The determination of the standard Gibbs energies of transfer and their importance for potential differences at the boundary between two immiscible electrolyte solutions are described in Sections 3.2.7 and 3.2.8. 

It is often more convenient to relate the potentiometer reading directly to concentration by adjusting the ionic strength and hence the activity of both the standards and samples to the same value with a large excess of an electrolyte solution which is inert as far as the electrode in use is concerned. Under these conditions the electrode potential is proportional to the concentration of the test ions. The use of such solutions, which are known as TISABs (total ionic strength adjustment buffers), also allows the control of pH and their composition has to be designed for each particular assay and the proportion of buffer to sample must be constant. 

Fig. 4-10. Electron energy levels in (a) an isolated solid metal and in (b) a metal electrode immersed in an electrolyte solution M = metal S = electrolyte solution e(STD) = gaseous electrons in the standard state e

Figures 4-11 and 4-12 show schematic energy diagrams for the electron transfer from the standard gaseous state through the electrolyte solution into the metal electrode. As mentioned in Chap. 2, the electron level (the real potential of electron) a s/v> in an electrolyte solution consists of an electrostatic energy...

Fig. 4-11. Energy diagram for electron transfer from a standard gaseous electron across a solution/vacuum interface, through an electrolyte solution, and across a metal/solution interface into a metal electrode = real potential of electrons e,s) in electrolyte...

If the sample solutions contain an indifferent electrolyte at a sufficiently high and constant concentration, then it can be assumed that the activity coefficients and Uquid-junction potentials are also constant and can be included in the value of the standard potential of the given ISE. It is then possible to calibrate the system using solutions with known concentrations of the ion sensed by the ISE and thus to obtain sample concentration values directly. 

The standard potential of these ISEs, analogous to halide ISEs, depends on the activities of silver and sulphur in the membrane. When metallic silver is used as a contact soldered directly onto the membrane, the silver activity in the membrane equals one and the Ag2 S ISE has properties identical with a silver electrode of the first kind or with a silver sulphide electrode of the second kind, depending on the solution with which it is in contact [203], The membrane that is in contact with electrolyte on both sides behaves similarly [106],... 

The electrochemical potential of the solution and semiconductor, see Fig. 3.6, are determined hy the standard redox potential of the electrolyte solution (or its equivalent the standard redox Fermi level, Ep,redo, and the semiconductor Fermi energy level. If these two levels do not lie at the same energy then movement of charge across the semiconductor - solution interface continues until the two phases equilibrate with a corresponding energy band bending, see Fig. 3.8. 

A bare surface of silicon can only exist in fluoride containing solutions. In reality, in these media, the electrode is considered to be passive due to the coverage by Si— terminal bonds. Nevertheless, the interface Si/HF electrolyte constitutes a basic example for the study of electrochemical processes at the Si electrode. In this system, the silicon must be considered both as a charge carrier reservoir in cathodic reactions, and as an electrochemical reactant under anodic polarization. Moreover, one must keep in mind that, according to the standard potential of the element, both anodic and cathodic charge transfers are involved simultaneously (corrosion process) in a wide range of potentials. 

The most stable oxidation states for protactinium are Pa(V) and Pa(IV). The chemical behavior of Pa(V) closely mimics that of Nb(V) and Ta(V), and experimental data are consistent with a 5f(l) rather than a 6d(l) electron configuration for the Pa(IV) species [37]. The electrochemical literature for Pa is mainly focused on the characteristics of the Pa(V)/Pa(IV) couple and the electrodeposition of Pa metal films from aqueous and nonaqueous electrolyte solutions. In aqueous solutions, only Pa(V) and Pa(IV) ions are known to exist, and the standard potential for the Pa(V)/Pa(IV) redox couple is in the range of —0.1 to -0.32 V [38]. 

The vast majority of electrochemical data on americium ions has heen obtained in aqueous solutions. Americium can exist in aqueous solutions in the oxidation states III, IV, V, and VI. The divalent state is difficult to attain in aqueous solutions because of the proximity of the standard potential of the Am(III)/Am(II) couple to the solvent/supporting electrolyte breakdown potential. Previous reviews have presented the formal and standard potentials for the various americium couples and these reviews should be consulted by the interested reader for more detailed discussion [133, 134]. Table 3 contains a summary of selected formal potentials Ef from these reviews in 1 M HCIO4 for convenience. AU values are calculated from various measurement techniques except for the Am(VI)/Am(V) couple (Am02 /Am02" "), which was determined directly. 

The analysis of the process of measuring potential differences across phase boundaries demonstrates the impossibility of using standard potential measuring devices to determine the value of a single metal-solution potential difference. The electrochemist proceeds, therefore, somewhat humbled but not defeated. He or she must ask how much information about the potential difference across an elec-trode/electrolyte interface can be obtained. 

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