Cathodic reaction area

The mechanism of poultice corrosion is shown diagramatically in Figure 4. Corrosion initially occurs uniformly over the whole metal surface. However, as the process continues, the oxygen reduction cathodic reaction may become restricted to a band near the surface where oxygen is readily available. Corrosion of metal then takes place preferentially slightly below the band. The pH rises in the cathodic reaction area due to an increase in hydroxyl ion concentration while the anodic area decreases in pH due to the hydrolysis of metal chloride reaction products. 

Reduction of oxygen is one of the predominant cathodic reactions contributing to corrosion. Awareness of the importance of the role of oxygen was developed in the 1920s (19). In classical drop experiments, the corrosion of iron or steel by drops of electrolytes was shown to depend on electrochemical action between the central relatively unaerated area, which becomes anodic and suffers attack, and the peripheral aerated portion, which becomes cathodic and remains unattacked. In 1945 the linear relationship between rate of iron corrosion and oxygen pressure from 0—2.5 MPa (0—25 atm) was shown (20). 

Typical platinum catalyst loadings needed to support the anodic and cathodic reactions are currently 1 to 2 mg/cm" oi active cell area. Owing to the cost of platinum, substantial efforts have been made to reduce the catalyst loading, and some fuel cells have operated at a catalyst loading of 0.25 mg/cm". 

Cathode—the electrode of an electrolytic cell where reduction takes place. During corrosion, this is the area at wliich metal ions do not enter the solution. During cathodic reactions, cations take up electrons and discharge them, hence reducing oxygen. That is, there is a reduction from a higlier to a lower state of valency. 

Since metals are electronic conductors, the anodic and cathodic reactions will not necessarily occur at the same site, and anodic and cathodic areas can develop as in aqueous solutions. For example, wash-line attack is often a feature of corrosion by fused salts in contact with air. 

Essentially, the pH is controlled to suppress the hydrogen evolution cathodic reaction- The Pourbaix Diagram for iron indicates that high pH values as well as low values may lead to corrosion. The construction of these diagrams for higher than ambient temperatures - shows how the area of the alkaline zones increases considerably under boiler conditions, so that the risk of corrosion is correspondingly higher. Many feed systems contain copper alloys. 

The corrosion current (it is also assumed that the area of the metal is 1 cm so that - ho ) occurs at a value within the Tafel region for the anodic and cathodic reaction, i.e. transport overpotential is negligible. 

When an electrode is at equilibrium the rate per unit area of the cathodic reaction equals that of the anodic reaction (the partial currents) and there is no net transfer of charge the potential of the electrode is the equilibrium potential and it is said to be unpolarised ... 

The oxides often are nonstoichiometric (with an excess or dehcit of oxygen). Many oxides are semiconducting, and their conductivity can be altered by adding various electron donors or acceptors. Relative to metals, the applications of oxide catalysts in electrochemistry are somewhat limited. Cathodic reactions might induce a partial or complete reduction of an oxide. For this reason, oxide catalysts are used predominantly (although not exclusively) for anodic reactions. In acidic solutions, many base-metal oxides are unstable and dissolve. Their main area of use, therefore, is in alkaline or neutral solutions. 

In cathodic area, the Tafel slope in the presence of DDTC is bigger than that in the absence of DDTC, and the cathodic curves imder the conditions of different DDTC concentration are almost parallel and their Tafel slopes only change a little. These demonstrate that the chemisorption of DDTC on the surface of jamesonite electrode also inhibits the cathodic reaction, but the chemisorption amoimt of DDTC is a little and almost not affected by the DDTC concentration due to their negatively electric properties of DDTC anion and the electrode surface. This reveals that there is a little DDTC chemisorption on the mineral even if the potential is lower (i.e., negative potential). 

Kendig and Leidheiser (16) electrochemica1ly evaluated thin (9 micron) polybutadiene coatings on steel. They concluded that movement of the corrosion potential in the noble direction was indicative of an increasing cathodic/anodic surface area ratio. Oxygen and water penetrate the coating to produce the cathodic reaction at the metaI/coat ng interface. 

Figures 16.8 and 16.9 show only the anodic polarization curves for corrosion cells. The important question is, where do these curves intersect with the polarization curves for likely cathodic reactions, such as hydrogen evolution or oxygen absorption The intersection point defines the corrosion current density icorr and hence the corrosion rate per unit surface area. As an example, let us consider the corrosion of titanium (which passivates at negative Eh) by aqueous acid. In Fig. 16.10, the polarization curves for H2 evolution on Ti and for the Ti/Ti3+ couple intersect in the active region of the Ti anode. To make the intersection occur in the passive region (as in Fig. 16.11), we must either move the H+/H2 polarization curve bodily...

In the previous analysis, homogeneous current distribution has been assumed but, on many occasions, corrosion occurs with localized attack, pitting, crevice, stress corrosion cracking, etc., due to heterogeneities at the electrode surface and failure of the passivating films to protect the metal. In these types of corrosion processes with very high local current densities in small areas of attack, anodic and cathodic reactions may occur in different areas of disparate dimensions. 

The basic mechanism for the instability of ultrapure metals was suggested by Wagner and Traud in a classic paper in 1938.1 The essence of their view is that for corrosion to occur, there need not exist spatially separated electron-sink and -source areas on the corroding metal. Hence, impurities or other heterogeneities on the surface are not essential for the occurrence of corrosion. The necessary and sufficient condition for corrosion is that the metal dissolution reaction and some electronation reaction proceed simultaneously at the metal/environment interface. For these two processes to take place simultaneously, it is necessary and sufficient that the corrosion potential be more positive than the equilibrium potential of the M, + + ne M reaction and more negative than the equilibrium potential of the electronation (cathodic) reaction A + ne — D involving electron acceptors contained in the electrolyte (Fig. 12.8). 

Consider a system consisting of a metal corroding in an electrolyte. The corrosion process involves a metal-dissolution deelectronation (anodic) reaction at electron-sink areas on the metal and an electronation (cathodic) reaction at electron-source areas. (This picture is applicable to a metal s corroding by a Wagner-Traud mechanism provided one imagines the sink and source areas shrunk to atomic-sized dimensions and considers the situation at one instant of time.)... 

Thus combining these expressions demonstrates that only in the cases in which the areas on which the anodic and cathodic reactions occur are equal can the anodic and cathodic current densities be equal ... 

This paradox of lowering both corrosion rates can be understood by considering the accounting of the electrons. By simply increasing the anode area, the number of electrons per second (i.e., the anodic current) that is produced is increased at all potentials more positive than the open circuit Econ of Metal 2. To accommodate this increased production, the cathodic reaction on Metal 1 must increase the rate at which it consumes electrons (i.e., the cathodic current). Increased cathodic current can only be achieved by moving the system to more negative potentials. The more negative the potential, the lower the dissolution rate of both materials. 

---