Metal-reaction equilibrium potential

FIGURE 26.35 Current-voltage Tafel plot showing cathodic protection by nse of a sacrificial anode. = equilibrium potential of the redaction reaction = equilibrium potential of the primary metal dissolution reaction E a = equilibrium potential of the sacrificial anode oxidation reaction. 

It is obvious that this reaction can only lead to metal dissolution, if the metal electrode potential is negative from the hydrogen electrode potential. This is the reason for the classification of metals into noble metals (the equilibrium potential is more positive than the standard hydrogen potential) and non-noble metals (the equilibrium potential is more negative than the standard hydrogen potential). The kinetic of the total process can be described by the Butler-Volmer equation for the two partial reactions. 

When two or more complementary processes such as those illustrated in Chap. 3 occur over a single metallic surface, the corrosion potential that results from such situations is a compromise between the various equilibrium potentials of all the anodic and cathodic reactions involved. The difference between the resultant potential (E) and each individual reaction equilibrium potential (E ) is called polarization and is quantified in terms of overpotential described in Eq. (5.1) ... 

Since any current resulting from tire anodic reaction must be consumed by tire catlrodic reaction, tire catlrodic current,7, must be equal to tire airodic current As a consequence, tire equilibrium potential of a metal (e.g. Fe) tlrat is immersed into air aqueous electrolyte will be adjusted by tire condition tlrat = j This is... 

Table 1.7 shows typical half reactions for the oxidation of a metal M in aqueous solutions with the formation of aquo cations, solid hydroxides or aquo anions. The equilibrium potential for each half reaction can be evaluated from the chemical potentials of the species involved see Appendix 20.2) and it should be noted that there is no difference thermodynamically between equations 2(a) and 2(b) nor between 3(a) and 3(b) when account is taken of the chemical potentials of the different species involved. 

Thus the tendency for an electrochemical reaction at a metal/solution interface to proceed in a given direction may be defined in terms of the relative values of the actual electrode potential E (experimentally determined and expressed with reference to the S.H.E.) and the reversible or equilibrium potential E, (calculated from E and the activities of the species involved in the equilibrium). 

It is apparent (Fig. 1.21) that at potentials removed from the equilibrium potential see equation 1.30) the rate of charge transfer of (a) silver cations from the metal to the solution (anodic reaction), (b) silver aquo cations from the solution to the metal (cathodic reaction) and (c) electrons through the metallic circuit from anode to cathode, are equal, so that any one may be used to evaluate the rates of the others. The rate is most conveniently determined from the rate of transfer of electrons in the metallic circuit (the current 1) by means of an ammeter, and if / is maintained constant it can eilso be used to eveduate the extent. A more precise method of determining the quantity of charge transferred is the coulometer, in which the extent of a single well-defined reaction is determined accurately, e.g. by the quantity of metal electrodeposited, by the volume of gas evolved, etc. The reaction Ag (aq.) -t- e = Ag is utilised in the silver coulometer, and provides one of the most accurate methods of determining the extent of charge transfer. 

Fig. 20.24 Potential energy-distance from metal surface curves, illustrating (a) an M /M system in which, owing to the relative position of the energy wells, the initial ionisation reaction occurs more rapidly than the discharge reaction, resulting in an excess negative charge on the surface of the metal, (b) equilibrium at which the energy wells are approximately the same and...

Potential-pH Equilibrium Diagram (Pourbaix Diagram) diagram of the equilibrium potentials of electrochemical reactions as a function of the pH of the solution. The diagram shows the phases that are thermodynamically stable when a metal reacts with water or an aqueous solution of specified ions. 

As shown in Fig. 33, the decreasing mechanism of this fluctuation is summarized as follows At a place on the electrode surface where metal dissolution happens to occur, the surface concentration of the metal ions simultaneously increases. Then the dissolved part continues to grow. Consequently, as the concentration gradient of the diffusion layer takes a negative value, the electrochemical potential component contributed by the concentration gradient increases. Here it should be noted that the electrochemical potential is composed of two components one comes from the concentration gradient and the other from the surface concentration. Then from the reaction equilibrium at the electrode surface, the electrochemical potential must be kept constant, so that the surface concentration component acts to compensate for the increment of the concen-... 

The areas bounded by solid lines correspond to regions of thermodynamic stability of certain substances that are named in the diagram. This stability is relative. The dashed line a in the diagram corresponds to the equilibrium potential of the hydrogen electrode. Metallic zinc, for which the reaction lines are below the line for the hydrogen electrode, can be oxidized while hydrogen is evolved (see Section 2.4.1). 

At most metals that are in contact with an electrolyte containing their own ions, the equilibrium potential of the metal s discharge and ionization reaction,... 

Some metals are thermodynamically unstable in aqueous solutions because their equilibrium potential is more negative than the potential of the reversible hydrogen electrode in the same solution. At such electrodes, anodic metal dissolution and cathodic hydrogen evolution can occur as coupled reactions, and their open-circuit potential (OCP) will be more positive than the equilibrium potential (see Section 13.7). 

In the mixed potential theory (MPT) model, both partial reactions occur randomly on the surface, both with respect to time and space. However, given the catalytic nature of the reductant oxidation reaction, it may be contended that such a reaction would tend to favor active sites on the surface, especially at the onset of deposition, and especially on an insulator surface catalyzed with Pd nuclei. Since each reaction strives to reach its own equilibrium potential and impose this on the surface, a situation is achieved in which a compromise potential, known as the mixed potential (.Emp), is assumed by the surface. Spiro [27] has argued the mixed potential should more correctly be termed the mixture potential , since it is the potential adopted by the complete electroless solution which comprises a mixture of reducing agent and metal ions, along with other constituents. However, the term mixed potential is deeply entrenched in the literature relating to several systems, not just electroless deposition. 

The simultaneous transfer of four electrons is unlikely, and the overall reaction must contain several steps. An important intermediate is hydrogen peroxide H2O2, and its occurrence makes it difficult to establish even the equilibrium potential experimentally. The reaction is further complicated by the fact that in aqueous solutions almost all metals are covered by an oxide film in the potential range over which the reduction occurs. 

The electromotive force of this cell gives the equilibrium potential Em2-/m for the transfer reaction of metal ions as shown in Eqn. 6-18 ... 

TABLE 6-2. Hie standard equilibrium potentials for transfer reactions of metal ions at metal electrodes at 25 C. [Handbooks of electrocfaemistiy.]... 

In electrochemistry, the chemical potential of hydrated ions has been determined from the equilibrium potential of ion transfer reactions referred to the normal hydrogen electrode. For the reaction of metal ion transfer (metal dissolution-deposition reaction) of Eqns. 6-16 and 6-17, the standard equilibriiun potential Sive in terms of the standard chemical potential, li, by Eqn. 

For metal electrodes, the anodic 03Q n reaction proceeds at electrode potentials more anodic than the equilibrium potential Bo of the reaction as shown in Fig. 10-14. For n-type semiconductor electrodes, the anodic photoexdted oxygen reaction proceeds at electrode potentials where the potential E of the valence band edge (predsely, the potential pEp of the quasi-Fermi level of interfadal holes, pCp = — CpEp) is more anodic than the equilibrium oxygen potential Eq, even i/the observed electrode potential E is less anodic than the equilibrium oxygen potential E03. The anodic hole transfer of the o Qgen reaction, hence, occurs at photoexdted n-type semiconductor electrodes even in the range of potential less anodic than the equilibriiun potential Eq of the reaction as shown in Fig. 10-14. 

When the cathodic reaction is the reduction of oi n molecules for which the equilibrium potential is relatively high (much more anodic than the corrosion potential), the corrosion current is frequently controlled by the diffusion of hydrated o Q en molecules towards the corroding metal electrode thus, the corrosion ciurent equals the diffusion current of o en molecules as shown in Fig. 11-8. For this mode of diffusion-controlled corrosion of metals the cathodic Tafel constant is... 

The mixed-potential model demonstrated the importance of electrode potential in flotation systems. The mixed potential or rest potential of an electrode provides information to determine the identity of the reactions that take place at the mineral surface and the rates of these processes. One approach is to compare the measured rest potential with equilibrium potential for various processes derived from thermodynamic data. Allison et al. (1971,1972) considered that a necessary condition for the electrochemical formation of dithiolate at the mineral surface is that the measmed mixed potential arising from the reduction of oxygen and the oxidation of this collector at the surface must be anodic to the equilibrium potential for the thio ion/dithiolate couple. They correlated the rest potential of a range of sulphide minerals in different thio-collector solutions with the products extracted from the surface as shown in Table 1.2 and 1.3. It can be seen from these Tables that only those minerals exhibiting rest potential in excess of the thio ion/disulphide couple formed dithiolate as a major reaction product. Those minerals which had a rest potential below this value formed the metal collector compoimds, except covellite on which dixanthogen was formed even though the measured rest potential was below the reversible potential. Allison et al. (1972) attributed the behavior to the decomposition of cupric xanthate. 

Table 4.1 shows the measured rest potential of sulphide electrode in thio collector solutions at pH = 6.86 and the equilibrium potential calculated for possible processes. In terms of the mixed potential model, the reaction products should be metal collector salts between four thio collectors and galena and jamesonite and should be disulphide between four thio collectors and pyrite and... 

---