Solvated-electron reference electrodes

Several of the key issues are reflected in the debate over the appropriate use of pe to describe redox conditions in natural waters (129-131). The parameter is defined in terms of the activity of solvated electrons in solution (i.e., pe = - log e ), but the species e aq does not exist under environmental conditions to any significant degree. The related concept of pe (132), referring to the activity of electrons in the electrode material, may have a more realistic physical basis with respect to electrode potentials, but it does not provide an improved basis for describing redox transformations in solution. The fundamental problem is that the mechanisms of oxidation and reduction under environmental conditions do not involve electron transfer from solution (or from electrode materials, except in a few remediation applications). Instead, these mechanisms involve reactions with specific oxidant or reductant molecules, and it is these species that define the half-reactions on which estimates of environmental redox reactions should be based. 

The transformation of metal-electrode surfaces via electrooxidation to their metallooxides, solvated ions, and metal complexes is fundamental to most anodic electrochemical processes (batteries, electrorefining, anodic-stripping analysis, and reference electrodes). Although this is traditionally represented as the removal of one (or more) electrons from a metal atom at the electrode... 

For reduction, relevant data from polarographic and cyclic voltammetric experiments are summarized in Tables 1 and 2, respectively. For the results in Table 1 the variety of solvents and reference electrodes used makes comparisons difficult. It is clear, however, that even with the activation of a phenyl substituent (entries 6,7,9-14) reduction occurs at very cathodic potentials. In this context it is worth noting that in aprotic solvents at ca. — 3 V vs. S.C.E.) it becomes difficult to distinguish between direct electron transfer to the alkyne and the production of the cathode of solvated electrons. Under the latter conditions the indirect electroreductions show many of the characteristics of dissolving metal reductions (see Section II.B). Even at extreme cathodic potentials it is not clear that an electron is added to the triple bond the e.s.r. spectra of the radical anions of dimesitylacetylene and (2,4,6,2, 4, 6 -hexa-r-butyldiphenyl)acetylene have been interpreted in terms of equal distribution of the odd electron in the aromatic rings . 

A special reference electrode is the electron electrode, which can operate in solvents capable of solvating electrons, such as ammonia. A sheet of platinum suspended in a solution of sodium metal in ammonia may be used as a reference electrode [208]. The concentration of the reference solution can be varied over relatively wide limits without a change in the potential of the electrode. A sodium concentration of 0.001 M was found suitable [208]. 

Ammonia has been employed mostly for cathodic reactions, but some oxidations [330,333] have been carried out in this medium, although the potential range in the anodic direction is quite small. The anodically limiting reaction is oxidation to nitrogen and protons [340] the cathodically limiting reaction is the transfer of electrons to the solvent, which occurs at about —2.3 V (versus Hg pool electrode) in a saturated solution of TBAI. In the elecltrolytic generation of solvated electrons the potential is determined by the surface concentration of electrons and no external reference electrode is needed. 

LiCl, NaN03, and tetraalkylammonium salts can be used as supporting electrolytes. For the electrolytic generation of solvated electrons mainly LiCl has been employed [343,344]. A reversible reference electrode in EDA is the Zn(Hg)/Zn" electrode [345], but the Hg pool [246] or the aqueous calomel electrode, fitted with a suitable salt bridge, is also applicable. 

The Birch and Benkeser reactions of some unsaturated organic compounds [318 and references therein], which consist of a reduction by sodium or lithium in amines, can be mimicked electrochemically in the presence of an alkali salt electrolyte. The cathodic reaction is not the deposition of alkali metal on the solid electrode but the formation of solvated electrons. Most of the reactions described were performed in ethylenediamine [319] or methylamine [308,320]. A feature of these studies is variety introduced by the use of a divided or undivided cell. In a divided cell, the product distribution appears to be the same as that in the classic reduction by metal under similar conditions. In contrast, in an undivided cell the corresponding ammonium salt is formed at the anode it plays the role of an in situ generated proton donor. Under such conditions, the proton concentration... 

Table 1 summarizes the basic relationships that link energy characteristics of excess electrons with the values measured by the aforementioned methods (see also Fig. 1). In the equations given therein, i.e. in Eqs. (5) and (6) w , w , and w denote respectively metal-to-vacuum, metal-to-solution, and solution-to-vacuum photoemission work functions AT is the Volta potential difference for a metal-solution system Eg is the equilibrium potential of the electrode in solvated electron solution and il(RE) is the Fermi level of the reference electrode. Equation (6) is approximate (see above) because the solvated electron entropy has not been taken into consideration. The main error in equating the heat of electron solvation and the activation energy of the thermoemission current for the solvated electron solution is caused by the variation in the solution s surface potential with temperature apparently, here specific adsorption of solvated electrons (or of an alkali metal) on the solution/vapour interface makes major contribution to the surface potential . This error can be probably neglected if measurements are taken in very dilute solutions (<10 mol/1, see ) of the alkali metal. This follows from the dependence measured in between thermoemission current and the concentration of sodium in hexamethylphosphotriamide. 

Now we shall draw the reader s attention to an interesting fact. As shown in Sect-tion 2.3, the difference between the delocalized and solvated electron levels for hexamethylphosphotriamide is by almost the same value, i.e. about 0.4 eV higher than for water or liquid ammonia, i.e. for the solvents having a branched structure of H bonds. It follows that the introduction of a hydrocarbon residue into the solvent s molecule forces out only the delocalized electrons from the polar medium the solvated electron energy level in all the enumerated solvents has almost the same value. (An independent confirmation to this is the closeness of equilibrium potentials of the electron in water, hexamethylphosphotriamide, and liquid ammonia — see Section 5 — vs. the reference electrode whose potential is independent of the solvent.)... 

Table 5 compares the standard potential of the electron electrode in hexamethylphosphotriamide (5 °C) with the standard potentials of alkali metals (25 °C). Data for liquid ammonia are also given. In both solvents the rubidium electrode potential serves as a reference point since it depends very little on the solvent. It is seen from the Table that in both solvents the standard equilibrium potential of the electron electrode is more positive than that of a lithium electrode and is close to the potentials of other alkali metals. In the course of experiment, cathodic production of dilute solutions (10 — 10 mol/1) of solvated electrons takes place and this makes the electron electrode equilibrium potential more positive compared to the standard value. In case of hexamethylphosphotriamide the same happens when electrons are bound in strong non-paramagnetic associates by the cations of all alkali metals except lithium (see Sect. 4). This enables one to assume that under the conditions of the experiments the electron-electrode equilibrium potential in liquid ammonia and hexamethylphosphotriamide is more positive than the equilibrium potential of all alkali metals. This makes thermodynamically possible primary cathodic generation of solvated electrons in solutions of all alkali metal salts in the two solvents. 

The splitting of redox reactions into two half cell reactions by introducing the symbol" e , which according to Eq.(II.28) is related to the standard electrode potential, is arbitrary, but useful (this e notation does not in any way refer to solvated electrons). When calculating the equihbrium composition of a chemical system, both e , and can be chosen as components and they can be treated numerically in a similar way equilibrium constants, mass balance, etc. may be defined for both. However, while represents the hydrated proton in aqueous solution, the above equations use only the activity of e , and never the concentration of e . Concentration to activity conversions (or activity coefficients) are never needed for the electron cf. Appendix B, Example B.3). 

When an inert metal (or any electronic conductor) is negatively polarized in the electrolyte (typically from 3 V versus the lithium reference electrode (LiRE) to 0 V versus LiRE), the following reactions take place at different potentials and at varying rates, depending on, on concentration, and on (q for ch of the following elec-trochemically active materials (1) solvents, (2) anions of the salts, and (3) impurities such as H2O, O2, HF, CO2 etc. (Fig. 1). Some of the products, especially at more positive potentials, may be soluble and diffuse away from the electrode, while others will precipitate on the surface of the electrode to form the SEI. At potentials lower than a few hundred millivolts, solvated electrons will be formed. These will also react with impurities, solvents and salts (e o scavengers) to produce similar products. The lifetime r, (r = 0.69A gfS] = rate constant ... 

The standard electrode potential (q.v.) is often not greatly different in non-aqueous solvents from that in water, although displacements due to differences in the strength of solvation of the ions are to be expected. The same reference electrodes as are used in water are also usually satisfactory. The rates of electrochemical reactions, however, can be radically altered by changes of solvent, since all the factors which govern the ease of transfer of electrons across the electrode surface are likely to be modified. These include the solvation of the electroactive ions, their tendency to ion-pairing and complex formation, the adsorbability of the solvent and of active species at the electrode surface, and the other factors that may affect the structure of the electrical double layer (q.v.). 

Electrochemical investigations have dealt with either U02(dik)2, or U(dik)4 complexes. Most investigations on U02(dik)2 derivatives have been carried out in coordinating solvents such as dmso or dmf, so that they actually refer to solvated molecules such as U02(dik)2(solvent)2 . Only in CHCI3 solution has it been assumed that the complex under study is actually U02(dik)2. In this medium, both U02(acac)2 and U02(dbm)2 undergo a chemically reversible U(VI) U(V) reduction. On the other hand, in coordinating solvents such as dmso or dmf, the one-electron reduction of most complexes is accompanied by chemical complications, which are probably due to the competitive coordination of the solvent itself (it is assumed that the diketonate may act as unidentate ligand in the uranium(V) species). The formal electrode potentials of the U(VI) U(V) process for selected complexes in various solvents are compiled in Table 18. 

The other interactions terms considered in the Hamiltoiuan of Eq. (12), such as spin and bonding between atoms, do not play any role in this case. In the course of the electron transfer the reactant changes its charge and hence its solvation this situation is illustrated in Fig. 10. In the adiabatic case the reactant shares its electron with the metal. We have referred the electronic energy to the Fermi level of the electrode, which is taken as zero for convenience. Thus, the occupation of the electronic state of the reactant is obtained by integrating the density of states up to this energy value (see Fig. 11). Since Delta is constant in this simple case, an analytical expression can be obtained "" ... 

According to electrochemical theory, the kinetics of an electrochemical reaction is controlled by the potential drop between the solid and solution phases [133-136]. A dynamic zone extending in both directions from the electrified interface over which this drop exists is called the double layer (DL) of charge. The DL in the solution is made up of adsorbed and solvated ions (molecules) and solvent. Its dense part, which is referred to as the Helmholtz layer (HL), plays the major role in the interfacial processes. At low ion concentration, there is also a diffuse layer Gouy layer) in the solution. The countercharged part of the DL in a metal electrode is comprised of a skin layer with an excess or a deficit of electrons. The DL in a semicondnctor electrode is called the space charge layer. It consists of an accumulation, depletion, or inversion layer with an excess or a deficit of electrons or holes and ionized donor or acceptor states, depending on... 

With reference to traditional electrolytes, the peculiarity of solid electrolytes results from their additional electronic conductivity. Other particularities may be mentioned the solidity of the electrolyte suppresses any convection, the ions are not solvated. Some solid electrolytes have a very simple crystalline structure where the energy profile of the path followed by the ions is known. This would make the situation ideally simple for electrode process studies. Unfortunately, this is frequently counteracted by the great complexi-ty of the electrode surface which is generally formed by a contact between two solids. This point has already been stressed by Raleigh. ... 

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