Corrosion of Metals

Corrosion and Faraday’s Laws

Corrosion

It is an electrochemical process where metals spontaneously oxidize to produce new substances by combining with oxygen or other non metals.
Some examples are rusting of iron, tarnishing of silver,formation of copper patina. Corrosion occurs when metal atoms lose electrons to become positive ions or get oxidized

Corrosion Prevention
Barrier Protection:

  • Painting the exterior
  • Electroplating or placing other metals on the exterior of the metal being protected
  • Greasing / Oiling

Cathodic Protection: Cause the Fe to be the cathode (In rusting of iron)

  • Galvanization: Using a more electropositive metal like zinc to be the anode and the metal to be protected usually Fe be the cathode and Zn is the anode. If the combination is Zn and Fe we say the iron is galvanized.
  • Sacrificial protection: The metal to be protected is made the cathode and a more reactive metal is connected by an external wire usually magnesium is the anode Example: Railway tracks, Ships, oil pipes connected to a chunk of magnesium using an external wire

Electroplating:

  • Electroplating is a method used to cover the surface of an object with a thin layer of metal. This is a process that can take place by placing a piece of metal to be coated on the cathode of a voltaic cell.

This can be explained using Faraday’s laws:

  • The relationship between electricity and electrochemical changes was first investigated by Michael Faraday in the 1830s.
  • He discovered that the mass of an element produced or consumed at an electrode was directly proportional to the time the celloperated, as long as the current was constant.
  • The charge of every mole of electrons that flows in the cell= 9.65 x 104C/mol