2.4 Corrosion processes: galvanic corrosion
When two dissimilar metals are in contact, or in close proximity with a conducting fluid in between, an electrochemical cell can be formed that leads to the more reactive metal becoming an anode and the less reactive metal a cathode.
This kind of corrosion is known as galvanic corrosion. It is not uncommon, since metals are often coated with others of different E0, and different metals are often in close contact with a common electrolyte.
One of the earliest examples of galvanic corrosion was recorded in the eighteenth century. The wooden hull of the Royal Navy frigate HMS Alarm (Figure 14) had been covered by copper sheathing, which was attached to the hull by iron nails.
One of the purposes of the copper sheath was to limit marine biofouling, which is known to plague many materials that are immersed in sea water. The growth of molluscs such as barnacles on the hulls of ships, which can then trap trailing seaweed, results in reduced speed and manoeuvrability. Copper limits fouling by inhibiting the attachment of molluscs. (Other organisms, such as bacteria, can also actually cause corrosion, as discussed in Box 2: Bacterial corrosion).
The hull was covered in 1761, and the copper sheath was found to be detached two years after fitting, during which time the Alarm had visited the Caribbean and elsewhere. The iron nails were found mostly to have corroded. Some nails remained intact, however, where their brown paper wrapping had remained in place between the copper and the iron, a fortuitous event that prevented total detachment of the sheath. The iron nails in contact with the copper were subject to rapid galvanic corrosion that led to detachment of the sheathing. The small anode (iron nails) to cathode (copper sheet) area ratio favoured the loss of the iron, as the rate of corrosion is directly proportional to the current density (a measure of electron flow). In a sense, the nails acted as local electron concentrators, so attack was rapid. Where it was present, the brown paper insulated the nails and so there was insufficient electron flow to cause corrosion.
The reason why marine environments are especially pernicious is the salt content of sea water. The presence of sodium and chloride ions increases the electrical conductivity compared with pure water, so galvanic or other cells formed between dissimilar metals react much faster.
Box 2: Bacterial corrosion
An unusual and perhaps unexpected corrosion problem can be caused by bacteria. As one of the oldest groups of organisms on the planet, bacteria have evolved to survive even in extreme environments. Bacterial corrosion can occur in fuel tanks, for example (Figure 15): fuel oil contaminates bilge water on tankers, and bacteria then grow profusely in the mixture.
The bacteria feed on the organic oil, releasing mild organic acids and depleting the oxygen content of the water. The acids will accelerate corrosion of the steel container, but a more serious stage can develop when certain species known as sulphate-reducing bacteria take over. These reduce the oxygen content of the sulphates commonly present in dirty fuel oils to produce hydrogen sulphide, or H2S. This compound is potent at corroding steel and can also enhance hydrogen embrittlement (which is a form of stress corrosion cracking), attack usually occurring as pits in the metal close to or under the bacterial colonies. Such colonies are perhaps better known for the ‘rusticles’ they produce – as were present on the wreck of the Titanic (Figure 16). The colonies of bacteria live on the rust, and promote further rusting through chemical attack of the underlying steel.
Such bacterial attack can also cause disasters directly, as in the gas explosion near Carlsbad in New Mexico, USA on 19 August 2000. The natural gas was carried in a 760 mm diameter steel pipe across a river via a suspension bridge. The pipe fractured suddenly, releasing gas that ignited into a fireball, engulfing the bridge and killing 12 people. It left a large crater, at the base of which were found the ends of the pipe; the missing pieces were ejected by the explosion (Figure 17).
Analysis of sludge found in the pipe showed evidence of extensive microbial attack in the form of deep pits in the pipe wall, and the presence of various contaminants including chlorides, hydrogen sulphide and sulphates. The fracture had occurred at a deeply corroded section of the 7.6 mm thick wall, where the wall thickness had been reduced to less than 2.5 mm. The rupture took the form of a 525 mm long crack along the axis of the pipe, which was under an internal pressure of 4.65 MPa. Better inspection procedures were recommended after the accident, including the use of cleaning ‘pigs’, which travel within pipes, both monitoring internal problems and cleaning debris away.
Calculate the approximate hoop stress in the pipe assuming no wall thinning, and then the effect of microbial corrosion on the hoop stress when the wall thickness has reduced to 2.5 mm.
You should recall that the hoop stress in a cylinder is given by:
where p is the internal pressure, r the radius of the cylinder and tc the wall thickness.The hoop stress acts such that the pipe will fail by a lengthways crack.
With the data provided, assuming no wall thinning, then:
With wall thinning to 2.5 mm due to the effects of corrosion, however, at failure:
Another example of a structure that was damaged by galvanic corrosion was the Statue of Liberty in New York harbour (Figure 18). Built in 1886 by Gustav Eiffel and Frederic Bartholdi, it was composed of an inner wrought-iron framework, with an outer cladding of copper attached by saddles of copper.
The risk of galvanic corrosion had been anticipated and so the two metals were separated by asbestos and shellac insulation. (Shellac is a natural resin that was widely used in the Victorian period as a lacquer or protective coating.) However, the shellac had degraded, and acidic rainwater had soaked the insulation, providing electrolytic conduction between the metals. The corrosion of the iron framework (Figure 19) was so extensive that there was concern it might collapse, and so in 1986 the statue was renovated. The wrought-iron framework was replaced by stainless steel, which will not corrode in the presence of copper, coated in a layer of PTFE insulation.
Highly localised attack, such as that found on the Statue, is also known as crevice corrosion, because attack is concentrated in the contact zone at the junction of the two dissimilar metals. A crevice forms and further attack occurs there, making the hole deeper (Figure 20). It is a common feature of corrosion, and can be contrasted with general overall attack. It is that much more dangerous since the damage is usually hidden from external inspection, until the strength of a product is reduced to a critical level and it fractures through the crevice. The loss of material lowers the section area, and there may also be a stress concentration within the crevice to magnify the stress further. Where the load levels are low, as in a galvanised water tank, nothing will happen until it leaks and alerts the owner to the problem. However, where a pressurised tank such as a boiler suffers the same problem, the effects may be much more dramatic.
Suggest why the rate of corrosion was lower on the Statue of Liberty than on HMS Alarm.
The Statue had corroded seriously in 100 years while HMS Alarm had corroded in only two years. The ship was subjected to continuous immersion in sea water, a good conductor owing to its high salt content, while the statue was subjected to only intermittent rain-water percolation through leaks in the outer copper skin. On the Alarm, the iron nails had a very small area, which meant that they corroded very quickly.
Explain the following observations of corrosion in terms of the electrochemical series:
(a) An empty, tin-plated steel food can will rust very rapidly after use if left outside.
(b) When the zinc coating on galvanized steel is broken, the underlying steel will rust only slowly.
(a) Tin (E0 = −0.14 V) is less reactive, with a less negative electrode potential, than iron (E0 = −0.44 V, Table 1). Iron will thus be attacked preferentially in the galvanic cells set up where the tin layer is broken. The rate of rusting will be rapid owing to the galvanic action set up between the two metals, especially in an external environment with exposure to slightly acidic rain.
(b) Zinc (E0 = −0.76 V) lies below iron (E0 = −0.44 V) with a more negative potential in the electrochemical series, so will corrode instead of the iron, which remains structurally intact until the zinc is consumed. After that, the iron will rust away until destroyed.