Electrochemistry

Corrosion

Corrosion is a natural oxidation-reduction reaction that occurs in metals. Atmospheric oxygen is the most common cathode in corrosion redox reactions.

Corrosion is the process by which metals oxidize and return from a reduced form to their natural oxidation state. Common examples of corrosion are the rusting of iron and the tarnishing of copper and silver.

The corrosion process is similar to the process that underlies galvanic cells. The metal acts as the anode because it loses electrons and is oxidized, forming metal ions. Atmospheric oxygen acts as the cathode. The electrons reduce atmospheric oxygen. The metal ions and the oxygen combine to form corroded metal. This process happens in the presence of an electrolyte, an ionic solution that can be decomposed by electricity, which is typically water.

Corrosion of iron in the presence of water involves multiple reactions. The iron metal loses electrons and forms metal ions in the water.
Fe(s)Fe2+(aq)+2e{\rm{Fe}}(s)\rightarrow{\rm{Fe}}^{2+}(aq)+2{\rm{e}}^-
The electrons reduce the atmospheric oxygen, forming hydroxide ions.
12O2(g)+H2O(l)+2e2OH(aq)\tfrac{1}{2}{\rm{O}}_2(g)+{\rm{H}}_2{\rm{O}}(l)+2{\rm{e}}^-\rightarrow2{\rm{OH}}^-(aq)
Iron ions and hydroxide ions react to form iron(II) hydroxide (Fe(OH)2).
Fe2+(aq)+2OH(aq)Fe(OH)2(s){\rm{Fe}}^{2+}(aq)+2{\rm{OH}}^-(aq)\rightarrow{\rm{Fe(OH)}}_2(s)
In the presence of oxygen, the iron(II) hydroxide (Fe(OH)2) forms iron rust (Fe2O3).
4Fe(OH)2(s)+O2(g)2Fe2O3H2O(s)+2H2O(l)4{\rm{Fe(OH)}}_2(s)+{\rm{O}}_2(g)\rightarrow2{\rm{Fe}}_2{\rm{O}}_3\!\cdot\!{\rm{H}}_2{\rm{O}}(s)+2{\rm{H}}_2{\rm{O}}(l)
Because the water sits on top of the metal, the metal ions can have one of two fates. They might go into solution in the water, leaving behind a pit in the metal surface. This allows more metal to become corroded in a process known as pitting. Alternatively, the metal ions can form a solid compound on the surface of the metal. This can form a barrier that prevents further corrosion of the metal surface. The path taken by the metal ion depends on the exact metal. Iron pits and rusts completely over time. Copper tarnish forms a barrier that prevents further contact with oxygen, stopping the corrosion.

Corrosion

Corrosion is a naturally occurring galvanic cell. The metal (in this case iron) acts as an anode, while atmospheric oxygen acts as a cathode in the presence of an electrolyte (usually water).
A great deal of energy and resources are put into forming metals for human use, so many methods to protect against corrosion have been developed. These methods generally rely on galvanic corrosion—corrosion between two different metals that favors one metal over the other. In galvanic corrosion, one of the metals acts as the cathode and reduces the ions in water to (OH) ions. This causes the half-reaction on the anode to proceed more quickly than it would under other circumstances. The metal that forms the cathode corrodes before the other metal corrodes. People have harnessed this phenomenon in a process called galvanization. Galvanization is the process that coats a metal with zinc as a sacrificial anode. A sacrificial anode is an anode made of a metal coupled to a more valuable metal, which it protects as part of a galvanic cell that undergoes galvanic corrosion. The zinc corrodes preferentially, protecting the metal underneath it. Thus, galvanization is a type of cathodic protection, a method of protecting a metal by sacrificing another metal as an anode in a galvanic cell.

Galvanization

Galvanization, a form of cathodic protection, uses zinc as a sacrificial anode. Zinc corrodes preferentially to iron when the two metals touch, slowing the corrosion of iron in favor of the corrosion of zinc.