Corrosion and rusting: experiments and protection
Using resources • Using materials
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Definition of rusting
Rusting describes the chemical corrosion of iron or steel, producing hydrated iron(III) oxide (commonly red-brown). The reaction requires oxidation of iron atoms, electron transfer to oxygen, and presence of an electrolyte to allow ion movement. Rusting weakens metal structures by converting strong metallic iron into brittle oxide products. Limiting factors include absence of oxygen, absence of liquid water, protective coatings, and use of less-reactive alloys. Salt water and acids increase the rate of rusting by enhancing electrical conductivity and promoting faster electron and ion movement.
Necessary conditions: air and water
Air supplies dissolved oxygen that accepts electrons during the oxidation of iron. Water acts as a medium for the movement of ions and often contains dissolved oxygen, salts, or acids that act as electrolytes. The combination of oxygen and an electrolyte allows the redox reactions that produce rust to occur. Cause → effect: presence of oxygen + presence of liquid water → electrochemical cells form on the metal surface → iron atoms oxidise to Fe2+ and Fe3+ → formation of hydrated iron(III) oxide (rust). Absence of either oxygen or water prevents the complete electrochemical cycle and so prevents rusting.
Classic experiments demonstrating requirements for rusting
Experiment setups commonly use clean iron nails placed in different environments: (1) open air and water, (2) sealed tube of dry air, (3) submerged under boiled water with a layer of oil to exclude air, and (4) in water with dissolved salt. Expected observations: rust forms where both air and water contact the nail; little or no rust forms where air or water is excluded. Interpretation of results uses controls and variable isolation. If a nail in boiled, oxygen-free water with oil shows no rust, the absence of oxygen prevents rusting despite water being present. If a nail in moist air but without liquid water shows reduced rusting, the lack of an electrolyte reduces the rate.
Sacrificial protection and relative reactivity
Sacrificial protection uses a more reactive metal attached to iron so that the more reactive metal oxidises preferentially. The more reactive metal donates electrons more readily and corrodes (sacrifices itself), protecting the iron cathodically. Cause → effect: more reactive metal present → oxidation occurs at the sacrificial metal surface → iron remains in the reduced state and does not oxidise → rusting of iron is prevented. Common sacrificial metals include zinc and magnesium because their positions in the reactivity series are above iron.
Other protection methods and limiting factors
Barrier methods prevent contact between iron and agents that cause corrosion. Paints, oil, plastic coatings, and galvanisation act as physical barriers. Galvanisation combines barrier and sacrificial effects when zinc coating is present: zinc delays corrosion by coating and then sacrifices if the coating is damaged. Limiting factors for protection effectiveness include coating continuity, presence of scratches, environmental chloride concentration, and galvanic coupling with other metals. Regular maintenance and correct choice of protection suit the environment to slow corrosion effectively.
Key notes
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