Corrosion and Electrochemical Corrosion Theory
Corrosion :
It is the gradual destruction of materials (usually metals) by chemical reaction with its environment. This means electrochemical oxidation of metals in reaction with an oxidant such as oxygen. Rusting, the formation of iron oxides, is a well-known example of electrochemical corrosion. This type of damage typically produces oxide(s) or salt(s) of the original metal. Corrosion degrades the useful properties of materials and structures including strength, appearance and permeability to liquids and gases.
Electrochemical Corrosion Theory:
Electrochemical corrosion involves two half-cell reactions; an oxidation reaction at the anode and a reduction reaction at the cathode. For iron corroding in water with a near neutral pH, these half cell reactions can be represented as:
Anode reaction: 2Fe => 2Fe2+ + 4e-
Cathode reaction: O2 + 2H2O + 4e- => 4OH-
There are obviously different anodic and cathodic reactions for different alloys exposed to various environments. These half cell reactions are thought to occur (at least initially) at microscopic anodes and cathodes covering a corroding surface. Macroscopic anodes and cathodes can develop as corrosion damage progresses with time.
From the above theory it should be apparent that there are four fundamental components in an electrochemical corrosion cell:
An anode.
A cathode.
A conducting environment for ionic movement (electrolyte).
An electrical connection between the anode and cathode for the flow of electron current. If any of the above components is missing or disabled, the electrochemical corrosion process will be stopped. Clearly, these elements are thus fundamentally important for corrosion control.
Importance of pH value to control the corrosion phenomenon :
Within the acid range (pH <4), the iron oxide film is continually dissolved. In cooling water, the potential for calcium carbonate precipitation increases with higher pH and alkalinity; thus the corrosion rate decreases slightly as pH is increased from 4 to 10. Above pH 10, iron becomes increasingly passive.
