Ionic electrode potential

The electrode potential can be defined not only by the energy level of electrons (the real potential ofelectrons)butalsoby theener gy/eue/oftons (the real potential of ions) in the electrode. The former maybe called the electronic electrode potenticd and the latter may be called the ionic electrode potential [Sato, 1995]. For instance, the electrode potential of a metal electrode can be defined in terms of the metal ion level (the real potential of metal ions), aM-ai/s/v), in the electrode as... 

For electrodes which have no electron energy levels in the energy range of general interest, such as ionic crystalline sohd electrodes and membrane electrodes, only the concept of ionic electrode potential can be of practical significance. 

In Chapter 4 the physical basis for the electrode potential is presented based on the electron and ion levels in the electrodes, and discussion is made on tiie electronic and ionic electrode potentials. Chapter 5 deals with the structure of... 

We may call the electrode potential defined by the ionic energy level the ionic electrode potential, and the electrode potential defined by the electronic energy level may be called the electronic electrode potential. In the case in which the electrode has no electronic level in the energy range of our interest such as certain membrane electrodes, it is convenient to describe the system in terms of the ionic electrode potential rather than the electronic electrode potential [Refs. 4 and 5.]. 

Stem layer adsorption was involved in the discussion of the effect of ions on f potentials (Section V-6), electrocapillary behavior (Section V-7), and electrode potentials (Section V-8) and enters into the effect of electrolytes on charged monolayers (Section XV-6). More speciflcally, this type of behavior occurs in the adsorption of electrolytes by ionic crystals. A large amount of wotk of this type has been done, partly because of the importance of such effects on the purity of precipitates of analytical interest and partly because of the role of such adsorption in coagulation and other colloid chemical processes. Early studies include those by Weiser [157], by Paneth, Hahn, and Fajans [158], and by Kolthoff and co-workers [159], A recent calorimetric study of proton adsorption by Lyklema and co-workers [160] supports a new thermodynamic analysis of double-layer formation. A recent example of this is found in a study... 

At low currents, the rate of change of die electrode potential with current is associated with the limiting rate of electron transfer across the phase boundary between the electronically conducting electrode and the ionically conducting solution, and is temied the electron transfer overpotential. The electron transfer rate at a given overpotential has been found to depend on the nature of the species participating in the reaction, and the properties of the electrolyte and the electrode itself (such as, for example, the chemical nature of the metal). 

In Section 8, the material on solubility constants has been doubled to 550 entries. Sections on proton transfer reactions, including some at various temperatures, formation constants of metal complexes with organic and inorganic ligands, buffer solutions of all types, reference electrodes, indicators, and electrode potentials are retained with some revisions. The material on conductances has been revised and expanded, particularly in the table on limiting equivalent ionic conductances. 

In addition to simple dissolution, ionic dissociation and solvolysis, two further classes of reaction are of pre-eminent importance in aqueous solution chemistry, namely acid-base reactions (p. 48) and oxidation-reduction reactions. In water, the oxygen atom is in its lowest oxidation state (—2). Standard reduction potentials (p. 435) of oxygen in acid and alkaline solution are listed in Table 14.10- and shown diagramatically in the scheme opposite. It is important to remember that if or OH appear in the electrode half-reaction, then the electrode potential will change markedly with the pH. Thus for the first reaction in Table 14.10 O2 -I-4H+ -I- 4e 2H2O, although E° = 1.229 V,... 

When two different metals are immersed in the same electrolyte solution they will usually exhibit different electrode potentials. If they are then connected by an electronic conductor there will be a tendency for the potentials of the two metals to move towards one another they are said to mutually polarise. The polarisation will be accompanied by a flow of ionic current through the solution from the more negative metal (the anode) to the more positive metal (the cathode), and electrons will be transferred through the conductor from the anode to the cathode. Thus the cathode will benefit from the supply of electrons, in that it will dissolve at a reduced rate. It is said to be cathodically protected . Conversely, in supplying electrons to the cathode the anode will be consumed more rapidly, and thus will act as a sacrificial anode. 

When metals are arranged in the order of their standard electrode potentials, the so-called electrochemical series of the metals is obtained. The greater the negative value of the potential, the greater is the tendency of the metal to pass into the ionic state. A metal will normally displace any other metal below it in the series from solutions of its salts. Thus magnesium, aluminium, zinc, or iron will displace copper from solutions of its salts lead will displace copper, mercury, or silver copper will displace silver. 

This procedure of using a single measurement of electrode potential to determine the concentration of an ionic species in solution is referred to as direct potentiometry. The electrode whose potential is dependent upon the concentration of the ion to be determined is termed the indicator electrode, and when, as in the case above, the ion to be determined is directly involved in the electrode reaction, we are said to be dealing with an electrode of the first kind . 

In the Nernst equation the term RT/nF involves known constants, and introducing the factor for converting natural logarithms to logarithms to base 10, the term has a value at a temperature of 25 °C of 0.0591 V when n is equal to 1. Hence, for an ion M+, a ten-fold change in ionic activity will alter the electrode potential by about 60 millivolts, whilst for an ion M2 +, a similar change in activity will alter the electrode potential by approximately 30 millivolts, and it follows that to achieve an accuracy of 1 per cent in the value determined for the ionic concentration by direct potentiometry, the electrode potential must be capable of measurement to within 0.26 mV for the ion M+, and to within 0.13 mV for the ion M2 +. ... 

By tradition, electrochemistry has been considered a branch of physical chemistry devoted to macroscopic models and theories. We measure macroscopic currents, electrodic potentials, consumed charges, conductivities, admittance, etc. All of these take place on a macroscopic scale and are the result of multiple molecular, atomic, or ionic events taking place at the electrode/electrolyte interface. Great efforts are being made by electrochemists to show that in a century where the most brilliant star of physical chemistry has been quantum chemistry, electrodes can be studied at an atomic level and elemental electron transfers measured.1 The problem is that elemental electrochemical steps and their kinetics and structural consequences cannot be extrapolated to macroscopic and industrial events without including the structure of the surface electrode. 

It is interesting to note, as pointed out to me by Mr. J. L. Hoard, that these considerations lead to an explanation of the stability of trivalent cobalt in electron-pair bond complexes as compared to ionic compounds. The formation of complexes does not change the equilibrium between bivalent and trivalent iron very much, as is seen from the electrode potentials, while a great change is produced in the equilibrium between bivalent and trivalent cobalt. 

A correlation between the amount of adsorbed ions and the electrode potential, in particular E. , has been identified apparently for the first time by Frumkin and Obrutschewa [26Fru]. A minimum of ionic adsorption was found at E, this is equivalent to the absence of specific adsorption at Ep c- The measurement of the amount of adsorbed ions was performed by measuring the ionic concentration in the solution as a function of the electrode potential or by measuring the surface concentration of adsorbed ions by e.g. radiotracer techniques (see also 4.2). (Data obtained with this method are labelled lA). 

Because of the interaction between the electrode surface and the various species present in the ionic phase an excess of charged species may be present on the ionically conducting side of the phase boundary. Its extent depends upon the actual electrode potential. At a certain electrode potential this excess vanishes, the corresponding potential value is called potential of zero charge. Knowledge of this... 

This is an example of a reversible reaction the standard electrode potential of the 2PS/PSSP + 2c couple is zero at pH 7. The oxidation kinetics are simple second-order and the presence of a radical intermediate (presumably PS-) was detected. Reaction occurs in the pH range 5 to 13 with a maximum rate at pH 6.2, and the activation energy above 22 °C is zero. The ionic strength dependence of 2 afforded a value for z Zg of 9 from the Bronsted relation... 

The physical concept of a single electrode potential has been also discussed in terms of the energy levels of ions in electrode systems. This concept may be usefirl in the cases where the system has no electronic energy levels in a range of practical interest, such as in ionic solid crystalline and electronically nonconductive membrane electrodes. "... 

All factors influencing the potentials of the inner or outer Helmholtz plane will also influence the zeta potential. For instance, when, owing to the adsorption of surface-active anions, a positively charged metal surface will, at constant value of electrode potential, be converted to a negatively charged surface (see Fig. 10.3, curve 2), the zeta potential will also become negative. The zeta potential is zero around the point of zero charge, where an ionic edl is absent. 

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