Corrosion of metals
Corrosion is the destruction of a metal by its environment. In chemical corrosion the metal reacts directly with an oxidiser (O₂, halogens, H₂S, SO₂); in electrochemical corrosion an electrolyte (damp, salty water) is present and two touching metals (or a metal and an impurity) form a galvanic couple in which the more active metal is destroyed. Iron rusts in damp air: Fe – 2e⁻ → Fe²⁺, O₂ + 2H₂O + 4e⁻ → 4OH⁻, and overall 4Fe + 6H₂O + 3O₂ → 4Fe(OH)₃ (rust, Fe₂O₃·nH₂O). Although aluminium is active, a dense Al₂O₃ film protects it. Methods of protection are paint, enamel and plastic coatings; metal coatings (zinc, chromium, nickel); adding inhibitors to the corrosive medium; making alloys (stainless steel); and electrochemical protection — a plate of a more active metal (a protector, e.g. Zn or Mg) is attached to the structure, becomes the anode and is destroyed while the iron, as the cathode, is saved. A zinc coating still protects iron when scratched (Zn is more active), but if a tin coating is scratched the iron rusts faster, because tin is less active than iron.
Observation (with an adult): three clear glasses hold (1) plain water, (2) salty water, (3) water covered with a layer of vegetable oil. Put one clean steel screw or small iron object in each, watch for a week and record the results. Do not use sharp iron objects (nails, needles); do the experiment with an adult. No acids are used.