Concept 07 / 07Electrochemistry
Corrosion: Electrochemical Mechanism of Rusting and its Prevention
Corrosion as an electrochemical process, the mechanism of rusting of iron with anodic and cathodic sites, conditions that accelerate corrosion, and methods of prevention including barrier coatings, galvanisation, sacrificial anodes, tin plating and alloying.
Corrosion is the slow deterioration of a metal by reaction with its environment: rusting of iron, tarnishing of silver, green coating on copper (basic copper carbonate). It causes enormous economic losses in buildings, bridges, ships and machinery.
Mechanism of rusting (see figure). A drop of water containing dissolved oxygen and carbon dioxide (which makes it slightly acidic) on an iron surface forms a miniature galvanic cell.
Anode (a region of the metal, often a strained or impure spot): , .
Electrons travel through the metal to a cathodic region, where oxygen is plentiful (near the edge of the drop):
Cathode: , .
Overall: , .
The ions are then oxidised by atmospheric oxygen to and deposit as hydrated ferric oxide, rust, . The ions migrate through the water film, which acts as the electrolyte.
Factors that accelerate corrosion: moisture and oxygen (both essential), electrolytes such as salt (sea water, road salt), acids and , contact with a less active metal (for example iron in contact with copper), impurities and strains in the metal, and higher temperature.
Prevention.
Barrier protection: covering the surface with paint, oil, grease, varnish or plastic, keeping out moisture and oxygen.
Galvanisation: coating iron with zinc. Zinc is more active (), so even if the coating is scratched, zinc is oxidised in preference to iron (sacrificial protection).
Sacrificial anode (cathodic protection): connecting a block of a more reactive metal (Mg or Zn) to underground pipelines, ship hulls or storage tanks; it corrodes instead of iron and is replaced periodically.
Tin plating (food cans): tin () is less reactive than iron and protects only while intact; if scratched, iron becomes the anode and corrodes faster than it would uncoated.
Alloying: stainless steel (with Cr and Ni) forms a thin protective oxide layer.
Anti-rust solutions and inhibitors, such as bisphenol or alkaline phosphate and chromate solutions, form protective films.
Key Takeaways:
- Rusting is electrochemical: iron is the anode, oxygen is reduced at the cathode.
- Water and oxygen are both required; salt and acids speed it up.
- Rust is .
- Zinc protects iron sacrificially (galvanising); tin protects only as a barrier.
- Pitfall: a scratched tin plate rusts faster than bare iron, while scratched galvanised iron is still protected.
Brain-Teaser 1: Why does iron rust faster near the sea and on roads treated with salt?
Solution: Dissolved salt makes the water film a much better electrolyte, so ions move more easily between the anodic and cathodic regions of the corrosion cell, increasing the current and hence the rate of iron dissolution.
Brain-Teaser 2: Explain why galvanised iron resists rusting even when scratched, whereas tinned iron does not.
Solution: Zinc has a more negative reduction potential than iron, so at a scratch, zinc becomes the anode and is oxidised while iron acts as the cathode and is protected. Tin has a less negative potential than iron, so at a scratch iron becomes the anode and corrodes, accelerated by contact with the tin cathode.
Brain-Teaser 3: Magnesium blocks are attached to a steel ship hull. If a block loses of magnesium in a year, what average current did it supply?
Solution: Moles of Mg , releasing of electrons: charge . One year , so the average current is .
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