Anodic E-i curves

Anodic E-i curves for nickel obtained by potentiostatic, potentiokinetic or, in earlier days, galvanostatic techniques, have been published by many workers. Unfortunately, good agreement is not always found between data from different sources. The principal reasons for the discrepancies appear to lie in the nature and amount of impurities in the metaP or in the solution -both of which may have a profound effect on the shape of the curve, and in variations in experimental procedure" . 

It is now appropriate to consider the kinetics of the anodic reaction with particular reference to the phenomenon of passivity, but since the mechanism is dealt with in detail in Section l.S this discussion will place the emphasis on the anodic E-i curves. 

Fig. 1.33 Regions of the potentiostatically determined anodic E-i curve for a metal that shows an active-passive transition...

A typical Evans diagrams for the corrosion of a single metal is illustrated in Fig. 1.26a (compare with Fig. 1.23 for two separable electrodes), and it can be seen that the E -I and E -I curves are drawn as straight lines that intersect at a point that defines and (it is assumed that the resistance for the solution is negligible). E can of course be determined by means of a reference electrode, but since the anodic and cathodic sites are inseparable direct determination of /co by means of an ammeter is not... 

Over the years the original Evans diagrams have been modified by various workers who have replaced the linear E-I curves by curves that provide a more fundamental representation of the electrode kinetics of the anodic and cathodic processes constituting a corrosion reaction (see Fig. 1.26). This has been possible partly by the application of electrochemical theory and partly by the development of newer experimental techniques. Thus the cathodic curve is plotted so that it shows whether activation-controlled charge transfer (equation 1.70) or mass transfer (equation 1.74) is rate determining. In addition, the potentiostat (see Section 20.2) has provided... 

Fig. 1.32 E-i curve for the simultaneous cathodic reduction of HjO (curve ImqpH) and dissolved oxygen (ABCG) which give the combined curve/4flC > E, Fis the anodic curve for...

Fig. 1.23 E-I curves for the corrosion of zinc (see Fig. 1.22) showing the relationship between Ef, p and i) for the cathodic and anodic half reactions...

Fig, 1.24 Tafel lines for a single exchange process. The following should be noted (a) linear f-log I curves are obtained only at overpotentials greater than 0-052 V (at less than 0-052 V E vs. i is linear) b) the extrapolated anodic and cathodic -log / curves intersect at tg the equilibrium exchange current density and (c) /, and the anodic and cathodic current densities... 

Skold and Larson" in studies of the corrosion of steel and cast iron in natural water found that a linear relationship existed between potential and the applied anodic and cathodic current densities, providing the values of the latter were low. However, the recognition of the importance of these observations is due to Stern and his co-workerswho used the term linear polarisation to describe the linearity of the rj — i curve in the region of E o , the corrosion potential. The slope of this linear curve, AE — AJ or Af - A/, is termed the polarisation resistance, / p, since it has dimensions of ohms, and this term is synonymous with linear polarisation in... 

It will be seen from the graphical representation that CIO- ions are discharged at a much lower anode potential, i. e. in preference to Cl- ions, if the concentration is the same for both kinds of ions. We can, therefore, conclude that in a concentrated chloride solution with a much lower hypochlorite content the oxidation of hypochlorite will occur simultaneously with the chlorine liberation at a potential necessary for the later process, as the current curve of the hypochlorite oxidation will be within the range of the discharge potentials of Cl ions even at very low concentrations of CIO" ions. 

Consider first the polarization curve (i.e., Tafel plot) for the anodic halfreaction occurring in corrosion of stainless steels (Fig. 16.8). The diagram for the active region is much the same as has been seen for other anodes (Figs. 15.4 to 15.7). As E i is increased to a certain specific value, however, a sudden and dramatic drop in the anodic current density i occurs, corresponding to formation of an oxide film. At higher E, i remains constant at a very low level (the horizontal scale in Fig. 16.8 is logarithmic), and the metal has become passive, that is, effectively immune from corrosion. 

The polarization curves in Fig. 4.16 permit an estimate of Icorr as the intersection of pairs of oxidation and reduction curves corresponding to the condition that Iox = Ired = Icorr. Actually, in this case, the corrosion is uniform, and the anodic and cathodic reactions are assumed to occur uniformly over the surface. Under this assumption, unit area is taken for analysis (A = 1 m2), and either E versus log i or E versus log I curves can be used in the analysis. However, the use of E versus log i curves obscures the fundamental basis on which corrosion rates are estimated... 

Figure 11.8.2 Experimental anodic stripping i-E curve for thallium. Experimental conditions 1.0 X 10 M T1+, 0.1 MKCl solution, = -0.7 V V5. SCE, = 5 min, V = 33.3 mV/s. Circles are theoretical points calculated from (11.8.2). [Reprinted with permission from I. Shain and J. Lewinson, Anal. Chem., 33, 187 (1961). Copyright 1961, American Chemical Society.]...

If the potential-current (E-i) characteristics of the individual reactions were measured, the reactions could be readily modeled as electrochemical reactions with the battery at open circuit as indicated by the processes in Figure 10. If dynamic electrode potential-current relationships were determined, the electrode is expected to show the classic Tafel slope behaviors as the exchange current of the anodic-cathodic equilibrium is shifted into either direction. From the Tafel curves a value for the Eq and Iq of the electrode could be defined. 

Upon polarization of either electrode, the cell potential moves along the oxidation and reduction curves as shown in Fig. 1.1. When the current through the cell is f, the potential of the copper and zinc electrodes is Cj cu and e zn > and each of the electrodes have been polarized by (Ceq.cu i.Cu) and (Ceq.zn i,z )- Upon further polarization, the anodic and cathodic curves intersect at a point where the external current is maximized. The measured output potential in a corroding system, often termed the mixed potential or the corrosion potential (Tcorr)> h the potential at the intersection of the anodic and the cathodic polarization curves. The value of the current at the corrosion potential is termed the corrosion current (Icon) and can be used to calculate corrosion rate. The corrosion current and the corrosion potential can be estimated from the kinetics of the individual redox reactions such as standard electrode potentials and exchange current densities for a specific system. Electrochemical kinetics of corrosion and solved case studies are discussed in Chapter 3. 

Now we have seen how we can determine the corrosion rate when the anode area and the cathode area are equally large. But it should be remembered that fliis is certainly not a general condition. Thus, generally it is more correct to use die basic rule, i.e. Ela = Eh Then die corrosion current and the corrosion potential are determined by means of E-log I diagram in the same way as we have shown by the E-log i diagram. As a basis for drawing the E-log I curve for each reaction, either the E-log i curve and die electrode area for each reaction, or experimentally recorded E-log I curves have to be known. We shall return to this, particularly in connection with galvanic elements (Section 7.3). 

---