Evans diagram corrosion rate

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). 

Fig. 1.27 Evans diagrams illustrating (a) cathodic control, (b) anodic control, (c) mixed control, (d) resistance control, (e) how a reaction with a higher thermodynamic tendency ( r, ii) may result in a smaller corrosion rate than one with a lower thermodynamic tendency and (/) how gives no indication of the corrosion rate...

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.

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... 

The less well known (so-called) Evans diagrams involve kinetics and allow one to make a rough first cut at the order of magnitude of a corrosion rate (for a pure surface without blocking oxide films). 

A better method, from the point of view of fundamentals, is to plot the log of the current densities of the anodic dissolution current and that of the cathodic partner reaction as a function of potential, but at a given pH, respectively. The common log i at which they intersect determines the corrosion rate. These Evans-Hoar diagrams are fundamentally correct and tell whether the corrosion will be significant. However, the relevant data, which would have to take into account the presence of oxide films, etc., is at present sparse, so that Evans-Hoar diagrams are largely of value for teaching principles and seldom for giving industrially useful information on demand. 

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. 

The corrosion rate of the iron can be directly predicted from the Evans diagram by considering two facts ... 

Figure 25 Evans diagram for Fe in acid showing use of conservation of charge to determine Eom and corrosion rate (7corr), given complete knowledge of the kinetic parameters involved.

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 ... 

In the presence of oxidizing species (such as dissolved oxygen), some metals and alloys spontaneously passivate and thus exhibit no active region in the polarization curve, as shown in Fig. 6. The oxidizer adds an additional cathodic reaction to the Evans diagram and causes the intersection of the total anodic and total cathodic lines to occur in the passive region (i.e., Ecmi is above Ew). The polarization curve shows none of the characteristics of an active-passive transition. The open circuit dissolution rate under these conditions is the passive current density, which is often on the order of 0.1 j.A/cm2 or less. The increased costs involved in using CRAs can be justified by their low dissolution rate under such oxidizing conditions. A comparison of dissolution rates for a material with the same anodic Tafel slope, E0, and i0 demonstrates a reduction in corrosion rate... 

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... 

In an environment with a constant redox condition (e.g., permanently aerated and/or constant pH), a condition not uncommon in industrial and environmental situations, corr could shift in the positive direction for a number of reasons. Incongruent dissolution of an alloy could lead to surface ennoblement. Alternatively, as corrosion progresses, the formation of a corrosion product deposit could polarize (i.e., increase the overpotential, i), for) the anodic reaction as illustrated in the Evans diagram of Fig. 4. Polarization in this manner may be due to the introduction of anodic concentration polarization in the deposit as the rate of transport of dissolved metal species away from the corroding surface becomes steadily inhibited by the thickening of the surface deposit i.e., the anodic half-reaction becomes transport controlled. 

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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