Corrosion as an Electrochemical Process

A piece of bare iron left outside where it is exposed to moisture will rust quickly. It will do so even more quickly if the moisture is salt water. The corrosion rate is enhanced by an electrochemical process in which a water droplet becomes a voltaic cell in contact with the metal, oxidizing the iron.

Considering the sketch of a water droplet (after Ebbing), the oxidizing iron supplies electrons at the edge of the droplet to reduce oxygen from the air. The iron surface inside the droplet acts as the anode for the process

The electrons can move through the metallic iron to the outside of the droplet where

Within the droplet, the hydroxide ions can move inward to react with the iron(II) ions moving from the oxidation region. Iron(II) hydroxide is precipitated.

Rust is then quickly produced by the oxidation of the precipitate.

The rusting of unprotected iron in the presence of air and water is then inevitable because it is driven by an electrochemical process. However, other electrochemical processes can offer some protection against corrosion. For magnesium rods can be used to protect underground steel pipes by a process called cathodic protection.

Cathodic Protection Against Corrosion

Underground steel pipes offer the strength to transport fluids at high pressures, but they are vulnerable to corrosion driven by electrochemical processes. A measure of protection can be offered by driving a magnesium rod into the ground near the pipe and providing an electrical connection to the pipe. Since the magnesium has a standard potential of -2.38 volts compared to -.41 volts for iron, it can act as a anode of a voltaic cell with the steel pipe acting as the cathode. With damp soil serving as the electrolyte, a small current can flow in the wire connected to the pipe. The magnesium rod will be eventually consumed by the reaction

while the steel pipe as the cathode will be protected by the reaction