Corrosion local cells

As the corrosion rate, inclusive of local-cell corrosion, of a metal is related to electrode potential, usually by means of the Tafel equation and, of course, Faraday s second law of electrolysis, a necessary precursor to corrosion rate calculation is the assessment of electrode potential distribution on each metal in a system. In the absence of significant concentration variations in the electrolyte, a condition certainly satisfied in most practical sea-water systems, the exact prediction of electrode potential distribution at a given time involves the solution of the Laplace equation for the electrostatic potential (P) in the electrolyte at the position given by the three spatial coordinates (x, y, z). 

Other hypotheses invoke corrosion in Grignard reactions by direct chemical action l.l()[. No evidence that distinguishes local-cell corrosion from direct chemical action is known to us. However, there is only weak evidence of the discrete Mg(l) intermediates that are required by the hypothesis of direct chemical action (Section 7.3.K). In addition, pitting is a characteristic ol local-cell corrosion, which may also be easier to integrate with the I) model that direct... 

The complete circuit in local-cell corrosion involves ion conduction in solution between the anodic and cathodic sites (Figure 7.39). When the solvent is very polar, e.g., water, ionic conduction is excellent and the anodic and cathodic sites can be macroscopically separated. The conductivities of solutions of RMgX and MgX io T>F7E. however, are very low. It is doubtful that the requirement for a complete electrical circuit will allow anodic and cathodic sites to be separated by much more than A-scale distances, if any, in DEB. In principle, anodic and cathodic sites, if they exist, inighi be detected by scanning electrochemical devices or STM. 

The implications of the surface area ratio SJ in Eq. (3.25) are particularly important in association with various forms of local cell corrosion such as pitting and stress corrosion cracking for which... 

Cathodic protection (overriding all local cell corrosion activity by applying a film of an active metal, and rendering the bulk uniformly cathodic). 

Copper-containing lead alloys undergo less corrosion in sulfuric acid or sulfate solutions than pure lead or other lead alloys. The uniformly dispersed copper particles give rise to local cells in which lead forms the anode and copper forms the cathode. Through this anodic corrosion of the lead, an insoluble film of lead sulfate forms on the surface of the lead, passivating it and preventing further corrosion. The film, if damaged, rapidly reforms. 

In most aqueous systems, the corrosion reaction is divided into an anodic portion and a cathodic portion, occurring simultaneously at discrete points on metallic surfaces. Flow of electricity from the anodic to the cathodic areas may be generated by local cells set up either on a single metallic surface (because of local point-to-point differences on the surface) or between dissimilar met s. 

Attack associated with nonuniformity of the aqueous environments at a surface is called concentration cell corrosion. Corrosion occurs when the environment near the metal surface differs from region to region. These differences create anodes and cathodes (regions differing in electrochemical potential). Local-action corrosion cells are established, and anodic areas lose metal by corrosion. Shielded areas are particularly susceptible to attack, as they often act as anodes (Fig. 2.1). Differences in concentration of dissolved ions such as hydrogen, oxygen, chloride, sulfate, etc. eventually develop between shielded and nearby regions. 

Shells, clams, wood fragments, and other biological materials can also produce concentration cell corrosion. Additionally, fragments can lodge in heat exchanger inlets, locally increasing turbulence and erosion-corrosion. If deposits are massive, turbulence, air separation, and associated erosion-corrosion can occur downstream (see Case History 11.5). 

Figure 1.62b shows the result of raising the potential of a corroding metal. As the potential is raised above B, the current/potential relationship is defined by the line BD, the continuation of the local cell anodic polarisation curve, AB. The corrosion rate of an anodically polarised metal can very seldom be related quantitatively by Faraday s law to the external current flowing, Instead, the measured corrosion rate will usually exceed... 

Fig. 1.62 Potential/current curves for a metal polarised (a) cathodically and (b) anodically. The horizontal intercepts xy, x y, x"y" with AB and CB represent the local cell currents respectively, and yz, y z, y z" the externally applied currents (cathodic and anodic). In bimetallic corrosion yz, y z, etc. will be the galvanic current /gjjv, flowing from to (see...

The driving force of a thermogalvanic corrosion cell is therefore the e.m.f. attributable to these four effects, but modified by anodic and cathodic polarisation of the metal electrodes as a result of local action corrosion processes. 

Very small amounts of copper taken into solution may cause considerable corrosion of more anodic metals elsewhere in the system, particularly zinc , aluminium , and sometimes steelSmall particles of copper deposited from solution set up local cells that cause rapid pitting. 

Fig. 19.16 Schematic E — I diagrams of local cell action on stainless steel in CUSO4 + H2SO4 solution showing the effect of metallic copper on corrosion rate. C and A are the open-circuit potentials of the local cathodic and anodic areas and / is the corrosion current. The electrode potentials of a platinised-platinum electrode and metallic copper immersed in the same solution as the stainless steel are indicated by arrows, (a) represents the corrosion of stainless steel in CUSO4 -I- H2 SO4, (b) the rate when copper is introduced into the acid, but is not in contact with the steel, and (c) the rate when copper is in contact with the stainless steel...

There are various different forms and mechanisms of localized concentration cell corrosion. Unfortunately, because one or more types may be manifest at any one time, the results can be quite devastating. 

Although each form of concentration cell may be considered a discrete form of corrosion, in practice, more than one type may occur simultaneously. These forms of corrosion are all characterized by localized differences in concentration of hydrogen, oxygen, chloride, sulfate, and other minerals, but especially oxygen (producing the so-called differential oxygen concentration cell, or differential-aeration cell). The basic mechanisms surrounding each of these specific forms of concentration cell corrosion are discussed next. 

Localized, concentration-cell corrosion (differential aeration corrosion), occurring as Tuberculation corrosion Crevice corrosion Under-deposit corrosion Pitting corrosion All forms of localized, concentration-cell corrosion are indirect attack type corrosion mechanisms. They result in severe metal wastage and can also induce other corrosion mechanisms, e.g. Stress corrosion Corrosion fatigue... 

The idea that metal corrosion could be due to local-cell action was put forward in 1830 by Auguste Arthur de la Rive, and became very popular. An extreme view derived from this idea is the assertion that perfectly pure metals lacking all foreign inclusions will not corrode. However, it does not correspond to reality. It was established long ago... 

The dissolution of zinc in a mineral acid is much faster when the zinc contains an admixture of copper. This is because the surface of the metal contains copper crystallites at which hydrogen evolution occurs with a much lower overpotential than at zinc (see Fig. 5.54C). The mixed potential is shifted to a more positive value, E mix, and the corrosion current increases. In this case the cathodic and anodic processes occur on separate surfaces. This phenomenon is termed corrosion of a chemically heterogeneous surface. In the solution an electric current flows between the cathodic and anodic domains which represent short-circuited electrodes of a galvanic cell. A. de la Rive assumed this to be the only kind of corrosion, calling these systems local cells. 

Fig. 11-1. Mixed electrode model (local cell model) for corrosion of metals i = anodic current for transfer of iron ions i = cathodic current of electron transfer for reduction of hydrogen ions.

Even single metals, however, are subject to aqueous corrosion by essentially the same electrochemical process as for bimetallic corrosion. The metal surface is virtually never completely uniform even if there is no preexisting oxide film, there will be lattice defects (Chapter 5), local concentrations of impurities, and, often, stress-induced imperfections or cracks, any of which could create a local region of abnormally high (or low) free energy that could serve as an anodic (or cathodic) spot. This electrochemical differentiation of the surface means that local galvanic corrosion cells will develop when the metal is immersed in water, especially aerated water. 

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