Evans diagram corrosion potential

Consider the two materials whose polarization curves are shown in Fig. 31. Both the polarization curves and the Evans lines are shown for both materials. Material 1 is the more noble material (i.e., it has a more positive Ec0II) and has a lower circuit corrosion rate when it is uncoupled. If the surface area of the two materials is the same and the materials are coupled, then the two material-solution interfaces must come to the same potential. In a manner identical to that used for the example of iron in acid used to introduce Evans diagrams, the potential and current at which this condition is met can be found by applying the conservation of charge to the sysytem ... 

The two dashed lines in the upper left hand corner of the Evans diagram represent the electrochemical potential vs electrochemical reaction rate (expressed as current density) for the oxidation and the reduction form of the hydrogen reaction. At point A the two are equal, ie, at equiUbrium, and the potential is therefore the equiUbrium potential, for the specific conditions involved. Note that the reaction kinetics are linear on these axes. The change in potential for each decade of log current density is referred to as the Tafel slope (12). Electrochemical reactions often exhibit this behavior and a common Tafel slope for the analysis of corrosion problems is 100 millivolts per decade of log current (1). A more detailed treatment of Tafel slopes can be found elsewhere (4,13,14). 

The sohd line in Figure 3 represents the potential vs the measured (or the appHed) current density. Measured or appHed current is the current actually measured in an external circuit ie, the amount of external current that must be appHed to the electrode in order to move the potential to each desired point. The corrosion potential and corrosion current density can also be deterrnined from the potential vs measured current behavior, which is referred to as polarization curve rather than an Evans diagram, by extrapolation of either or both the anodic or cathodic portion of the curve. This latter procedure does not require specific knowledge of the equiHbrium potentials, exchange current densities, and Tafel slope values of the specific reactions involved. Thus Evans diagrams, constmcted from information contained in the Hterature, and polarization curves, generated by experimentation, can be used to predict and analyze uniform and other forms of corrosion. Further treatment of these subjects can be found elsewhere (1—3,6,18). 

The equilibrium potentials and E, can be calculated from the standard electrode potentials of the H /Hj and M/M " " equilibria taking into account the pH and although the pH may be determined an arbitrary value must be used for the activity of metal ions, and 0 1 = 1 is not unreasonable when the metal is corroding actively, since it is the activity in the diffusion layer rather than that in the bulk solution that is significant. From these data it is possible to construct an Evans diagram for the corrosion of a single metal in an acid solution, and a similar approach may be adopted when dissolved O2 or another oxidant is the cathode reactant. 

Evans Diagram diagram in which the E vs. I relationships for the cathodic and anodic reactions of a corrosion reaction are drawn as straight lines intersecting at the corrosion potential, thus indicating the corrosion current associated with the reaction. 

Fig. 12.17. The Evans diagrams are plots of the potentials of the two reactions (a) vs. the magnitude of the two currents or (b) vs. their logarithms. The intersections of the curves define the corrosion current and corrosion potential.

The above two methods ofpreventing corrosion can be understood easily with an Evans diagram (Fig. 12.39 see Section 12.19). (These diagrams it will be recalled, result from the superposition of the potential-current curves of the electronation and... 

With the help of the Evans diagram (Fig. E12.1) explain the influence of an inhibitor on the corrosion potential and corrosion current. Assume that the inhibitor decreases the exchange current density for the cathodic reaction. (Contractor)... 

It is well known that Pourbaix diagrams give the thermodynamic limits of corrosion. However, it is possible that corrosion in a system may be limited by kinetics to rates so low that corrosion that is thermodynamically possible can be neglected under practical circumstances. In this light, (a) construct an Evans diagram, i.e., a plot of the actual relevant electrode potentials against log i for... 

Consider an Evans diagram in a general way. The anodic dissolution reaction is to be represented in the Tafel region the same applies for the cathodic partner reaction, (a) Draw the two Tafel lines and show the region of intersection Oanodic= Cathodic)- Indicate on the graph the corrosion rate and corrosion potential. 

Figure 24 Schematic Evans diagram and polarization curve illustrating the origin of the negative hysteresis observed upon cyclic polarization for materials that do not pit. Line a represents the (unchanging) cathodic Evans line. Line b represents the anodic Evans line during the anodically directed polarization, while line c represents the anodic Evans line for the material after its passive film has thickened because of the anodic polarization. The higher corrosion potential observed for the return scan (E (back)) is due to the slowing of the anodic dissolution kinetics.

Figure 27 Schematic (a) Evans diagram and (b) corrosion potential vs. time behavior for localized corrosion stabilization. Line a on the Evans diagram represents the electrochemical behavior of the material before localized corrosion initiates, while line b represents the electrochemical behavior of the material in the localized corrosion site. Due to the low Tafel slope of the active site, the corrosion potential of the passive surface/local-ized corrosion site falls. If repassivation occurs, the anodic behavior reverts back to line a, and the corrosion potential increases again (line c). If repassivation does not occur, the corrosion potential will remain low (line d).

The information required to predict electrochemical reaction rates (i.e., experimentally determined by Evans diagrams, electrochemical impedance, etc.) depends upon whether the reaction is controlled by the rate of charge transfer or by mass transport. Charge transfer controlled processes are usually not affected by solution velocity or agitation. On the other hand, mass transport controlled processes are strongly influenced by the solution velocity and agitation. The influence of fluid velocity on corrosion rates and/or the rates of electrochemical reactions is complex. To understand these effects requires an understanding of mixed potential theory in combination with hydrodynamic concepts. 

The understanding gained by considering the Evans diagrams allows us to measure the corrosion current in a straightforward manner. First we must realize that the corrosion potential is in fact the open-circuit potential of a system undergoing corrosion. It represents steady state, but not equilibrium. It resembles the reversible potential in that it can be very stable. Following a small perturbation, the system will return to the open-circuit corrosion potential just as it returns to the reversible potential. It differs from the equilibrium potential in that it does not follow the Nemst equation for any redox couple and there is both a net oxidation of one species and a net reduction of another. 

Further oxidation results in the formation of hydrated ferric oxide or Fe(ni) hydroxide, i.e. rust. The corrosion potential (Ec) and corrosion current (/c) for the cathodic and anodic reaction can be represented by an Evans-type polarisation diagram, Eig. 6.6. Corrosion inhibitors interfere with the anodic or the cathodic partial reaction, or with both, resulting in a reduction in the corrosion current. 

These concepts can be applied quite directly to the corrosion behavior of iron. The effect of diffusion control on the corrosion rate was shown in the Evans Diagrams (b) and (c) of Figure 4. In the cathodic case the current becomes constant over a wide range of potentials because of control by the rate of diffusion of the reactant oxygen gas to the cathode. The anodic reaction can be dependent on the rate of diffusion of Fe++ from the anode. The rates in both cases can be Increased by stirring. The rate of Fe" Ion removal from the anode can also be increased by the presence of complexlng agents in the solution. 

The Evans diagram ( ) is a graphical presentation in semilogarithmic coordinates of the anodic and cathodic reaction rates expressed as partial currents dependent on potential. The basis for the Evans diagram is the corrosion model discussed above ... 

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