Prevention of Rusting

• A critical negative feature of rust is that it is not a protective layer; it is highly porous and flaky.
• Because it easily flakes off, fresh iron is continually exposed to oxygen and water, meaning the metal will continue to rust until it is completely destroyed.

Proving the Conditions for Rusting

• How can we prove beyond doubt exactly what makes iron rust?
• Chemists use a standard three-test-tube experiment involving iron nails to demonstrate that both oxygen and water are essential.
• Tube 1 (Control): An iron nail is placed in regular water and left open to the air. Both reactants are present, so the nail rusts heavily.
• Tube 2 (Excluding Oxygen): Water is boiled first to drive out any dissolved oxygen, and a layer of oil is poured over the surface. The oil prevents atmospheric oxygen from re-dissolving into the water, resulting in no rust.
• Tube 3 (Excluding Water): A nail is placed in a sealed tube containing a desiccant, such as anhydrous calcium chloride. The desiccant absorbs all moisture from the trapped air, resulting in no rust.

Prevention via Barrier Methods

• The simplest way to protect iron is to physically block the environment from reaching the metal surface.
• A barrier method works by providing a physical coating that prevents contact between the iron and the oxygen and water in the surroundings.
• Different barriers are chosen based on the object's function:
  • Painting: Ideal for large, stationary structures like bridges, ships, and fences.
  • Oiling or Greasing: Essential for moving parts, such as bicycle chains or engine components, where a brittle coat like paint would simply crack and flake off.
  • Plastic Coating: Often used for wire mesh or dish racks.
  • Electroplating: Using electrolysis to apply a thin layer of an unreactive metal (like silver or chromium) over the iron.
• The main drawback of barrier methods is that they do not offer protection once the coating is compromised. If the paint or plastic is scratched, the exposed iron will rust, and the porous rust can spread underneath the remaining barrier.

Prevention via Sacrificial Protection

• Imagine having a bodyguard that willingly takes a hit so you do not have to.
• Sacrificial protection involves attaching blocks of a metal that sits higher in the reactivity series (such as zinc or magnesium) directly to the iron.
• Because the sacrificial metal is more reactive, it loses its outer electrons much more easily than iron does.
• The more reactive metal is oxidised in preference to the iron, effectively "sacrificing" itself.

$$Zn(s) \rightarrow Zn^{2+}(aq) + 2e^-$$

• The electrons lost by the zinc flow into the iron, which actively prevents the iron atoms from losing electrons to become $Fe^{3+}$ ions.
• Unlike barrier methods, sacrificial protection continues to work perfectly even if the metal is heavily scratched, provided the two metals remain in electrical contact.
• Note that using a less reactive metal, such as copper, does not work. If copper were attached to iron, the iron would oxidise first to protect the copper, actually speeding up the rusting process.

Galvanising: The Dual Method

• You have probably seen shiny, mottled grey dustbins or watering cans — these have been treated with a special dual-action defence called galvanising.
• Galvanising involves coating an iron or steel object with a layer of zinc, usually by dipping it into molten zinc or through electroplating.
• It protects the iron in two distinct ways:
  1. Barrier Action: The zinc coating physically prevents contact between the iron and oxygen/water. The zinc reacts with air to form a tough, non-porous layer of zinc carbonate ($ZnCO_3$).
  2. Sacrificial Action: If the zinc coating becomes scratched and the iron is exposed, the iron still does not rust because the zinc acts as a sacrificial anode, oxidising in preference to the iron.