1.6 - Rusting

Corrosion of Metals

• Corrosion is a redox process in which a metal is oxidised spontaneously when its atoms release electrons to form metal ions: $\mathrm{M}\to {\mathrm{M}}^{\mathrm{n}+}+\mathrm{n}{\mathrm{e}}^{-}$.
• Rusting is the corrosion of iron. Both water and oxygen are required; salts and acids increase the rate by improving electrical conduction.
• Aluminium, zinc, lead, nickel and chromium form compact oxide layers that protect the underlying metal. Rust is porous and flakes off, continually exposing fresh iron.

Experiment 1D — corrosion of copper and iron

• Aim: Study corrosion of copper and iron.
• Problem statement: Do copper and iron undergo corrosion?
• Hypothesis: Copper corrosion produces a blue precipitate; iron corrosion produces a brown precipitate.
• Variables: type of metal; presence and colour of precipitate; salt solution kept constant.
• Materials: $20\mathrm{cm}$ copper wire, $20\mathrm{cm}$ iron wire and the salt/alkali solutions below.
• Apparatus: Test tubes, test-tube rack, sandpaper, beaker, tripod, wire gauze and Bunsen burner.
• Procedure:
  1. Clean the wires with sandpaper, coil them and place in separate test tubes.
  2. Mix $12\mathrm{c}{\mathrm{m}}^3$ of $0.5\mathrm{mol}\mathrm{d}{\mathrm{m}}^{-3}$ sodium chloride, $6\mathrm{c}{\mathrm{m}}^3$ of $0.05\mathrm{mol}\mathrm{d}{\mathrm{m}}^{-3}$ potassium chloride and $2\mathrm{c}{\mathrm{m}}^3$ of $0.5\mathrm{mol}\mathrm{d}{\mathrm{m}}^{-3}$ sodium hydroxide; heat the mixture.
  3. Cover both wires with the mixture and leave for three days.
• Results: Copper produces a blue precipitate; iron produces a brown precipitate. Both metals undergo corrosion.
• Metal atoms are oxidised and dissolved oxygen is reduced to $\mathrm{O}{\mathrm{H}}^{-}$. Copper ions form blue copper(II) hydroxide; iron corrosion ultimately gives brown hydrated iron(III) oxide.

Rusting as a Redox Reaction

• Different regions on wet iron act as tiny electrodes: the centre of a water droplet has less oxygen and forms the anode; the edge has more oxygen and forms the cathode.

• At the anodic region, iron is oxidised: $\mathrm{Fe}(\mathrm{s})\to \mathrm{F}{\mathrm{e}}^{2+}(\mathrm{aq})+2{\mathrm{e}}^{-}$

• Electrons move through the iron to a cathodic region, where oxygen is reduced: ${\mathrm{O}}_2(\mathrm{g})+2{\mathrm{H}}_2\mathrm{O}(\mathrm{l})+4{\mathrm{e}}^{-}\to 4O{\mathrm{H}}^{-}(\mathrm{aq})$

• $\mathrm{F}{\mathrm{e}}^{2+}$ and $\mathrm{O}{\mathrm{H}}^{-}$ form iron(II) hydroxide, which is further oxidised to hydrated iron(III) oxide, $\mathrm{F}{\mathrm{e}}_2{\mathrm{O}}_3\cdot \mathrm{x}{\mathrm{H}}_2\mathrm{O}$, or rust.

• Formation of iron(II) hydroxide: $\mathrm{F}{\mathrm{e}}^{2+}+2O{\mathrm{H}}^{-}\to \mathrm{Fe}(\mathrm{OH}{)}_2$. In rust, $x$ has variable integer values.

• Iron is the reducing agent; oxygen is the oxidising agent.

Effect of Contact with Other Metals

• When iron contacts a more electropositive metal such as Mg, Al or Zn, that metal releases electrons and is oxidised in preference to iron. Rusting of iron slows or stops.
• When iron contacts a less electropositive metal such as Sn, Pb or Cu, iron becomes the anodic metal, loses electrons more readily and rusts faster.

Experiment 1E — effects of other metals

• Problem: How do different metals in contact with iron affect rusting?
• Hypothesis: A more electropositive metal protects iron; a less electropositive metal accelerates rusting.
• Aim: Study the effects of other metals on rusting.
• Variables: Manipulated = metal coiled around iron; responding = rusting of iron; fixed = iron nail.
• Materials: Agar, phenolphthalein, potassium hexacyanoferrate(III), iron nails, magnesium ribbon, tin, zinc and copper strips.
• Apparatus: Test tubes, rack and sandpaper.
• Procedure:
  1. Clean five iron nails and the four metals with sandpaper. Coil one nail with each metal; leave the fifth nail uncoiled as the control.
  2. Place the nails in separate test tubes and completely immerse them in hot agar containing potassium hexacyanoferrate(III) and phenolphthalein.
  3. Leave for one day and examine the colours.
• Precaution: Potassium hexacyanoferrate(III) solution is poisonous.
• Blue colour indicates $\mathrm{F}{\mathrm{e}}^{2+}$ formation; pink indicates $\mathrm{O}{\mathrm{H}}^{-}$ formation. Colour locations identify oxidation and reduction regions.

Preventing Rusting

• Barrier protection isolates iron from water and oxygen: paint, oil or grease, and plastic coating. Damage to the coating exposes iron.

• Coating with another metal:

  • galvanising with zinc protects by a barrier and by sacrificial action even if scratched;
  • tin plating is a barrier, but scratched tin exposes iron beside a less electropositive metal and can accelerate rusting.

• Chrome plating protects items such as car bumpers, bicycles and decorative objects.

• Sacrificial protection: connect iron or steel structures to a more electropositive metal such as magnesium or zinc. The sacrificial metal corrodes instead and must be replaced periodically.

• Sacrificial protection is used on bridge pillars, ship hulls and underground pipes; magnesium protects iron pipes and zinc blocks protect ship hulls.

• Alloying: stainless steel contains iron alloyed with carbon, chromium and nickel. Its protective oxide layer resists corrosion; it is used for cutlery and surgical instruments.