Two-metal corrosion

The dry cell battery is a typical example of galvanic corrosion, or two metal corrosion as it is otherwise called. When two dissimilar metals are immersed in a conductive or corrosive medium, there is always the potential for a change in them. Once these metals are connected this difference induces electron flow between them. The less corrosion resistant metal is attacked more than the more resistant metal. This is an electrochemical process. In the case of a dry cell battery, the carbon electrode acts as the cathode (the more resistant materials) and zinc as the corroding anode. The natural phenomenon of corrosion is used in this case for producing electricity. 

While ordinary galvanic corrosion (two-metal corrosion) can be identified by a visual inspection alone, for thermogalvanic corrosion we need also to know a little about the service conditions. 

A potential (emf) difference usually exists between two dissimilar metals when they are immersed in a corrosive or conductive solution. If these metals are placed in contact (or electrically connected), this potential difference produces an electron flow between them. Corrosion of the less corrosion-resistant metal is usually increased and attack of the more resistant material is decreased, as compared with the behavior of these metals when they are not in contact. The less-resistant metal becomes anodic and the more-resistant metal cathodic. Usually the cathode or cathode metal corrodes very little or not at all in this type of couple. Because of the electric currents and dissimilar metals involved, this form of corrosion is called galvanic or two-metal corrosion. 

Galvanic or two-metal corrosion can influence erosion-corrosion when dissimilar metals are in contact in a flowing system. 

Corrosion problems are particularly important when two metals are in contact. The more reactive metal becomes the cathode of the cell and goes into solution when the cell is activated by an electrolyte. A typical cell is shown in Figure 13.7. When the metal in contact with iron is more reactive than iron itself, the iron is protected from corrosion. This is important when mechanical strength... 

Another ak pollutant that can have very serious effects is hydrogen sulfide, which is largely responsible for the tarnishing of silver, but also has played a destmctive role in the discoloration of the natural patinas on ancient bronzes through the formation of copper sulfide. Moreover, a special vulnerabihty is created when two metals are in contact. The electromotive force can result in an accelerated corrosion, eg, in bronzes having kon mounting pins. 

Galvanic Corrosion. Galvanic corrosion occurs when two dissimilar metals are in contact in a solution. The contact must be good enough to conduct electricity, and both metals must be exposed to the solution. The driving force for galvanic corrosion is the electric potential difference that develops between two metals. This difference increases as the distance between the metals in the galvanic series increases. 

Table 4 shows a galvanic series for some commercial metals and alloys. When two metals from the series are in contact in solution, the corrosion rate of the more active (anodic) metal increases and the corrosion rate of the more noble (cathodic) metal decreases. 

Generai description. Galvanic corrosion refers to the preferential corrosion of the more reactive member of a two-metal pair when the metals are in electrical contact in the presence of a conductive fluid (see Chap. 16, Galvanic Corrosion ). The corrosion potential difference, the magnitude of which depends on the metal-pair combination and the nature of the fluid, drives a corrosion reaction that simultaneously causes the less-noble pair member to corrode and the more-noble pair member to become even more noble. The galvanic series for various metals in sea water is shown in Chap. 16, Table 16.1. Galvanic potentials may vary with temperature, time, flow velocity, and composition of the fluid. 

Most galvanic corrosion processes are sensitive to the relatively exposed areas of the noble (cathode) and active (anode) metals. The corrosion rate of the active metal is proportional to the area of exposed noble metal divided by the area of exposed active metal. A favorable area ratio (large anode, small cathode) can permit the coupling of dissimilar metals. An unfavorable area ratio (large cathode, small anode) of the same two metals in the same environment can be costly. 

The attraction of rubbed amber and some other effects of electricity were known in ancient times. We know from finding nails in an old wreck that the Romans knew about contact corrosion combined with electric current flow. A skin of lead as a protection against boring worms covered the wooden planks of the ship and was nailed down with copper nails. Galvanic coupIe.s formed between the lead and the copper nails and the less noble lead sheets around the nails corroded in the seawater and fell off. The shipbuilders discovered a simple solution and covered the heads of the copper nails with lead as well. Galvanic current flow between the two metals was eliminated and corrosion was prevented (26). 

Galvanic corrosion is the enhanced corrosion of one metal by contact with a more noble metal. The two metals require only being in electrical contact with each other and exposing to the same electrolyte environment. By virtue of the potential difference that exists between the two metals, a current flows between them, as in the case of copper and zinc in a Daniell cell. This current dissolves the more reactive metal (zinc in this case), simultaneously reducing the corrosion rate of the less reactive metal. This principle is exploited in the cathodic protection (Section 53.7.2) of steel structures by the sacrificial loss of aluminum or zinc anodes. 

The corrosion potentials of the two metals in the environment under consideration will determine the direction of the transfer of electrons, but will provide no information on the rate of electron transfer, i.e. the magnitude of the galvanic current. Thus if E an.. is more positive than corr..B thc transfer of electrons will be from to with a consequent increase in the corrosion potential (more positive) of and a decrease in that of A/ the corrosion rate of will consequently increase and the corrosion rate of A/ will decrease compared with the rates when the metals... 

Even discounting the case of aluminium, which is usually covered by protective oxide films, it is evident from Table 1.23 that the quantitative connection between the galvanic corrosion rate of the more active member of the couple and the difference of reversible potentials of the two metals, is non-existent. 

Fig. 1.69 Effect of resistivity of solution on the distribution of corrosion on the more negative metal of a bimetallic couple, (a) Solution of very low resistivity and (b) solution of very high resistivity. Note that when the resisitivity is high the effective areas of the cathodic and anodic metals are confined to the interface between the two metals...

The British Non-Ferrous Metals Research Association carried out two series of tests, the results of which have been given by Gilbert and Gilbert and Porter these are summarised in Table 4.12. In the first series tough pitch copper tubes were exposed at seven sites for periods of up to 10 years. The two most corrosive soils were a wet acid peat (pH 4-2) and a moist acid clay (pH 4-6). In these two soils there was no evidence that the rate of corrosion was decreasing with duration of exposure. In the second series phosphorus-deoxidised copper tube and sheet was exposed at five sites for five years. Severe corrosion occurred only in cinders (pH 7 1). In these tests sulphides were found in the corrosion products on some specimens and the presence of sulphate-reducing bacteria at some sites was proved. It is not clear, however, to what extent the activity of these bacteria is a factor accelerating corrosion of copper. 

Zirconium, like titanium, depends upon the integrity of a surface film, usually of oxide, for its corrosion resistance, but there are differences in behaviour between the two metals when they are exposed to aggressive aqueous environments. 

Both metals are applied to copper-base alloys, stainless steels and titanium to stop bimetallic corrosion at contacts between these metals and aluminium and magnesium alloys, and their application to non-stainless steel can serve this purpose as well as protecting the steel. In spite of their different potentials, zinc and cadmium appear to be equally effective for this purpose, even for contacts with magnesium alloys Choice between the two metals will therefore be made on the other grounds previously discussed. 

There is no intermetallic compound formation and the electrodeposit behaves as a simple mixture of the two metals. It can be considered as basically a stable wick of tin through which zinc is fed to be consumed at a rate lower than its consumption from a wholly zinc surface. If the conditions are such that zinc is rapidly consumed, and no protective layer of corrosion products is formed, the coating may break down, but in mildly corrosive conditions some of the benefits of a zinc coating, without some of its disadvantages, are obtained. 

When two metals in intimate contact are subjected to vibration, a dark powder forms at the areas of contact. The effect is referred to as fretting corrosion though it is due to wear rather than true corrosive attack. The galling effect between nickel and steel ensures good resistance to fretting corrosion and lubricated nickel against steel is a very satisfactory combination used widely in industry for components assembled by press-fitting. 

The explicit aims of boiler and feed-water treatment are to minimise corrosion, deposit formation, and carryover of boiler water solutes in steam. Corrosion control is sought primarily by adjustment of the pH and dissolved oxygen concentrations. Thus, the cathodic half-cell reactions of the two common corrosion processes are hindered. The pH is brought to a compromise value, usually just above 9 (at 25°C), so that the tendency for metal dissolution is at a practical minimum for both steel and copper alloys. Similarly, by the removal of dissolved oxygen, by a combination of mechanical and chemical means, the scope for the reduction of oxygen to hydroxyl is severely constrained. 

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