Corrosion as a Battery: Evans Diagrams and Tafel Kinetics

Corrosion is a battery you did not build on purpose

Metal corrosion is an electrochemical reaction: at anodic sites the metal gives up electrons and dissolves as ions (oxidation), and at cathodic sites those electrons are consumed by a reduction reaction — usually oxygen reduction in aerated water, or hydrogen evolution in acid. Both half-reactions run on the same piece of metal, wired together by the metal itself as the electron conductor and the electrolyte film as the ionic path, which is exactly the layout of a battery, just one nobody wanted.

Reading an Evans diagram

Plot the anodic current density (dissolution) and the cathodic current density (reduction) against electrode potential on the same log-current axis and you get an Evans diagram. Each curve individually follows Butler–Volmer kinetics, but far from equilibrium each simplifies to a straight line on a semi-log plot — a Tafel line — with slope set by the Tafel constants βa (anodic) and βc (cathodic):

i = i_corr * [ exp((E - E_corr)/beta_a) - exp(-(E - E_corr)/beta_c) ]

  i_corr = corrosion current density  (sets the corrosion rate)
  E_corr = corrosion (mixed) potential, where anodic and cathodic
           currents are equal in magnitude and opposite in sign

Where the two Tafel lines cross is the corrosion potential E_corr and the corrosion current i_corr — the single point where total anodic current equals total cathodic current, so the net measurable current at that potential is zero even though metal is actively dissolving underneath. Faraday's law then converts i_corr directly into a mass-loss rate, which is why a two-line log-log sketch is enough to estimate how many microns per year a bare steel plate will lose in seawater.

What moves the lines: oxygen, pH, area ratio

Anything that raises the cathodic reaction rate shifts its Tafel line up and pushes i_corr higher — more dissolved oxygen, more agitation feeding oxygen to the surface, a more efficient cathode. Anything that passivates the anode — a stable oxide film, a chromium-rich stainless surface, a coating — bends the anodic line over and can collapse i_corr by orders of magnitude, which is exactly how stainless steel and aluminium survive despite being thermodynamically eager to oxidise. The relative area of anode to cathode matters just as much as the chemistry: a small anodic scratch on a large cathodic surface (a chipped tin can, a scratched galvanised sheet with the zinc breached) concentrates all the anodic current into a tiny spot and corrodes it very fast — the classic mechanism behind pitting.

Passivation and pitting

Many engineering metals are protected not by inertness but by a thin, self-healing oxide layer a few nanometres thick — passivation. On the Evans diagram this shows up as a region where the anodic current drops sharply and stays low across a wide potential range. Chloride ions are notorious for locally breaking that film, reinitiating rapid anodic dissolution in a tiny spot while the surrounding passive area keeps acting as a large, efficient cathode — the geometry that makes pitting corrosion so much more dangerous per gram of metal lost than uniform corrosion.

Fighting it: cathodic protection and coatings

If corrosion is a battery, you can rig the battery in your favour. Sacrificial anodes — zinc or magnesium blocks bolted to a ship's hull or a buried pipeline — are more anodic than steel, so they dissolve preferentially and hold the protected structure at a potential where its own anodic current is negligible. Impressed current systems do the same thing with an external power supply instead of a consumable metal. Barrier coatings simply try to keep the electrolyte off the metal in the first place, which is why a single scratch through paint on a car is worse than no paint at all in that spot: it creates a tiny anode surrounded by a huge cathodic area under the intact coating.

Frequently asked questions

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