Evans diagram solution

Figure 8.4. Current-potential curves for the reduction of Cu ions and the oxidation of reducing agent Red, formaldehyde, combined into one graph (an Evans diagram). Solution for the Tafel line for the reduction of Cu ions O.IM CUSO4, 0.175M EDTA, pH 12.50, Egq (Cu/Cu ) = -0.47 V versus SCE for the oxidation of formaldehyde 0.05 M HCHO and 0.075 M EDTA, pH 12.50, (HCHO) = -1.0 V versus SCE temperature 25 0.5°C. (From Ref. 10, with permission from the American Electroplaters and Surface Finishers Society.)...

Fig. 3. Hypothetical Evans diagram and polarization curve for a metal corroding in an acidic solution, where point A represents the current density, /q, for the hydrogen electrode at equiUbrium point B, the exchange current density at the reversible or equiUbrium potential, for M + 2e and point...

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... 

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. 

Figures 1.27a to d show how the Evans diagram can be used to illustrate how the rate may be controlled by either the polarisation of one or both of the partial reactions (cathodic, anodic or mixed control) constituting corrosion reaction, or by the resistivity of the solution or films on the metal surface (resistance control). Figures 1. lie and/illustrate how kinetic factors may be more significant than the thermodynamic tendency ( , u) and how provides no information on the corrosion rate.

In an Evans diagram 89> the mixed potential can easily be found and also be verified by measuring the open circuit potential of a zinc-amalgam electrode in a Cu2+-ion solution. Even the complication by the simultaneous presence of another reducible species, e.g., Pbz+ can be graphically demonstrated for different limiting conditions... 

Fig. 5. Tentative mixed potential model for the sodium-potassium pump in biological membranes the vertical lines symbolyze the surface of the ATP-ase and at the same time the ordinate of the virtual current-voltage curves on either side resulting in different Evans-diagrams. The scale of the absolute potential difference between the ATP-ase and the solution phase is indicated in the upper left comer of the figure. On each side of the enzyme a mixed potential (= circle) between Na+, K+ and also other ions (i.e. Ca2+ ) is established, resulting in a transmembrane potential of around — 60 mV. This number is not essential it is also possible that this value is established by a passive diffusion of mainly K+-ions out of the cell at a different location. This would mean that the electric field across the cell-membranes is not uniformly distributed.

Figure 8. Wagner-Traud (Evans) diagram for aluminum in aqueous solution, in the absence of dissolved oxygen.

Fig. 2. Current-potential curves in Evans diagram [29] format for reduction of Cu2+ ions and oxidation of H2CO. and are the equilibrium, or open circuit, potentials for the Cu2+ reduction and H2CO oxidation reactions, respectively. Assuming negligible interfering reactions, the vertical dashed lines indicate the exchange current densities for the two half reactions, and the deposition current for the complete electroless solution. Adapted from ref. 23.

The oxidation rate of a stainless steel is studied in an aqueous solution ofpH 3.1 saturated with H2 and free of 02. The following data were obtained from an Evans diagram at 25 °C. 

Return to the Fe dissolution experiment discussed above, altering the solution to contain 5 pM M Fe2+. In addition, allow hydrogen evolution to occur on the iron surface with an exchange current density of 1(T5 A/cm2, whereas the exchange current density for the iron reaction is 1CT6 A/cm2. Assume that both reactions have Tafel slopes of 100 mV/decade. These conditions are illustrated graphically in the Evans diagram, named in honor of its creator, U. R. Evans, shown in Fig. 25. The lines represent the reaction kinetics of the two reactions considered. 

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 ... 

Figure 5 Schematic Evans diagram and resulting potential-controlled polarization curve for a material that undergoes an active-passive transition and is in a reducing solution. The heavy line represents the applied currents required to polarize the sample.

Figure 6 Schematic Evans diagram and resulting potential-controlled polarization curve for a material that undergoes an active-passive transition and is in an oxidizing solution. The heavy line represents the applied currents required to polarize the sample. If the sample did not undergo an active-passive transition, it would corrode at a much higher rate in this solution, as is indicated by the intersection of the dotted line and the cathodic curve.

Figure 28 Schematic Evans diagrams and polarization curves for a material in a solution containing a redox couple that acts as a chemical potentiostat. The i used in the Evans diagram for the O/R redox couple is that relevant to the material of interest. In the absence of the redox couple, the material obtains Ec, i. In the presence of the redox couple, the material obtains Econ2. If Econ2 is above the pitting potential, the material will be rapidly attacked.

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. 

Figure 2 Evans diagram illustrating the influence of solution velocity on corrosion rate for a cathodic reaction under mixed charge transfer-mass transport control. The anodic reaction shown is charge transfer controlled.

A preliminary knowledge of which reaction steps could be key in determining the overall corrosion rate can be assessed by measurements of Corr as a function of important system parameters, e.g., oxidant concentration, solution composition, temperature. The proximity of ACOrr to either eM/Mn+ or /Red can indicate which of the two half-reactions may be rate determining. This is illustrated in Fig. 3A, which shows an Evans diagram for the combination of a fast anodic reaction coupled to a slow cathodic one. The corrosion of iron or carbon steel in aerated neutral solution would be an example of such a combination. The anodic reaction requires only a small overpotential (1) = /Mn+ - Ecorr) to sustain the corrosion current, /COrr, compared to the much larger overpotential required to sustain the cathodic reaction at this current. The anodic reaction would... 

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