The earliest experiments on cathodic protection were performed with zinc anodes that had been electrically connected to copper plates immersed in seawater. As can be seen on the galvanic series, such an arrangement would produce a cathode (copper) and an anode (zinc). In the large galvanic cell so formed, the zinc cylinder corroded away in a manner to protect the copper substrate. This method of cathodic protection can be used with other combination of metals providing the necessary current to the metal to be protected, as Sir Humphry Davy and Michael Faraday illustrated almost two centuries ago.
When two metals are electrically connected to each other in a electrolyte e.g. seawater, electrons will flow from the more active metal to the other, due to the difference in the electrical potential, the so called 'driving force'. When the most active metal (anode) supplies current, it will gradually dissolve into ions in the electrolyte, and at the same time produce electrons, which the least active (cathode) will receive through the metallic connection with the anode. The result is that the cathode will be negatively polarized and hence be protected against corrosion. To calculate the rates at which these processes occur, one has to understand the electrochemical kinetics associated with the complex sets of reactions that can all happen simultaneously on these metals.
