Corrosion

 Corrosion reactions are redox reactions in which the elemental form of a metal is oxidized to produce an unwanted compound or a water soluble ion

• Oxidized forms of a metal may not have the strength and conductive properties of the metal in the elemental form

Diatomic oxygen in the atmosphere has a relatively high reduction potential - higher than many metals

• For metals exposed to air, oxygen (in the presence of H⁺) is reduced to H₂O, and the metal is oxidized, in a spontaneous redox reaction

For some metals, however, the initial oxidation on the surface protects the underlying metal against further oxidation. This is because the metal oxide layer is impervious to both oxygen and H₂O (a source of H⁺)

• This is the case for aluminum, which although quite reactive upon exposure to oxygen and water, develops a metal oxide surface layer that protects the metal underneath

The corrosion of iron (i.e. rusting) costs the U.S. economy about $70 billion annually

• Formation of rust requires both O₂ and H₂O
• Anode reaction:

Fe(s) ® Fe²⁺(aq) + 2e^- E⁰_red = -0.44 V

• Cathode reaction:

O₂(g) + 4H⁺(aq) + 4e^- ® 2H₂O(l) E⁰_red = 1.23 V

• The Fe²⁺ that is formed eventually is oxidized further to Fe³⁺ which forms a hydrated iron (III) oxide (i.e. rust)

If iron is painted, then oxygen and water are prevented from contacting the metal, and corrosion is avoided

Sometimes the iron is coated with a thin layer of another metal, such as Tin (the layer of Tin keeps oxygen and water away from the iron)

• If the layer of Tin is worn away or scratched, the presence of Tin will actually accelerate the corrosion of the underlying iron:

Sn²⁺(aq) + 2e^- ® Sn(s) E⁰_red = -0.14 V

Fe²⁺(aq) + 2e^- ® Fe(s) E⁰_red = -0.44 V

• Iron will serve as the anode, and Tin the cathode in a redox reaction that will see the Iron oxidized (and actually oxygen will be reduced at the Tin cathode)

Another metal used to coat iron is Zinc, in a producing galvanized iron

• The Zinc will keep oxygen and water away from the iron
• However, if the Zinc coating is worn or scratched, the Zinc coating will electrochemically protect the Iron from oxidation:

Zn²⁺(aq) + 2e^- ® Zn(s) E⁰_red = -0.76 V

Fe²⁺(aq) + 2e^- ® Fe(s) E⁰_red = -0.44 V

• In a redox reaction involving Iron and Zinc, the Zinc will serve as the anode, and Iron the cathode.
• The zinc anode will oxidize and provide electrons for the reduction of Fe²⁺(aq) to elemental iron!
• This is called Cathodic Protection. The Zinc anode is termed a Sacrificial Anode.
• Iron pipes buried in the ground, and designed to carry water, would normally be expected to rust pretty quickly
• If they are buried along with a piece of Zinc, and connected by a wire, the zinc will provide cathodic protection (Magnesium will also work)

• This is also why aluminum-hulled boats will have a small piece of magnesium connected by a wire to the hull and in contact with the water. The magnesium will preferentially oxidize and drive the reduction of the aluminum hull. If the magnesium were not there, or once all the magnesium is oxidized, the aluminum will oxidize and the boat's hull will start to dissolve!