Questions: Corrosion Chemistry and Protection

Question 1

Stainless steel resists corrosion because it contains at least 10.5% chromium. What is the mechanism of this protection?

• A. Chromium makes the steel harder, physically preventing chemical attack
• B. Chromium displaces iron at the surface, and chromium is thermodynamically noble (unreactive)
• C. Chromium spontaneously forms a thin (1-5 nm), dense, self-healing Cr2O3 passive film that acts as a kinetic barrier to further oxidation, dramatically slowing the anodic dissolution rate
• D. Chromium increases the melting point of the steel, making it resistant to thermal oxidation

Answer: (C)

Chromium is actually thermodynamically more reactive than iron — it has a more negative standard reduction potential. However, chromium oxide (Cr2O3) is extremely dense, adherent, and chemically stable, forming a passive film only 1-5 nm thick that reduces the corrosion rate by orders of magnitude. This film is self-healing: if scratched or damaged, chromium in the underlying alloy re-oxidizes to repair the film almost instantly in oxygen-containing environments. The 10.5% threshold is the minimum chromium content needed to form a continuous passive film. This is a kinetic, not thermodynamic, form of protection — the steel wants to corrode but cannot because the passive film blocks ion transport.

Question 2

Pitting corrosion is more dangerous than uniform corrosion despite removing less total metal, because pits concentrate material loss into small areas that can penetrate through the wall thickness of pipes, tanks, and structural members, causing leaks or sudden failure with little visible warning.

• T. True
• F. False

Answer: True

Uniform corrosion spreads material loss evenly over the surface and is easily monitored by thickness measurements — a corrosion allowance can be designed into the component. Pitting corrosion, by contrast, creates localized attack at specific sites (often where the passive film has been disrupted by chloride ions, inclusions, or mechanical damage). The pit interior becomes anodic (metal dissolution) while the surrounding passive surface acts as a large cathode, creating an unfavorable area ratio that accelerates local attack. Pits can penetrate several millimeters deep while the surrounding surface appears pristine. The autocatalytic nature of pitting (dissolved metal ions hydrolyze, lowering pH inside the pit, which further destabilizes the passive film) means that once initiated, pits tend to grow aggressively. This makes pitting the most common cause of unexpected failure in stainless steels and aluminum alloys in chloride-containing environments.

Question 3

A zinc coating on steel (galvanizing) protects the steel even after the coating is scratched and the steel is exposed. Explain why, using electrochemical principles.

Model answer: Zinc is more electrochemically active (more negative reduction potential, -0.76 V vs SHE) than iron (-0.44 V vs SHE). When both metals are exposed to an electrolyte, they form a galvanic couple in which zinc acts as the anode (preferentially dissolving) and steel acts as the cathode (protected from dissolution). The zinc sacrificially corrodes to protect the steel — this is cathodic protection. The protection extends to exposed steel near the zinc boundary because the galvanic current flows through the electrolyte, polarizing the steel cathodically. This is why galvanized steel retains corrosion resistance even with scratches, cuts, or minor damage to the coating, unlike a paint or polymer barrier coating where any breach exposes the underlying steel to direct attack.

Sacrificial protection is one of the most important principles in corrosion engineering. The galvanic series — the ranking of metals and alloys by their corrosion potential in a given environment — predicts which metal in a couple will corrode preferentially. Zinc, magnesium, and aluminum alloys are all commonly used as sacrificial anodes to protect steel structures (ship hulls, pipelines, offshore platforms). The alternative approach, impressed-current cathodic protection, uses an external power supply to force the protected structure cathodic, achieving the same electrochemical effect without consuming a sacrificial metal.

Question 4

Why does the corrosion rate of iron increase dramatically in the presence of dissolved oxygen, even though oxygen does not directly attack the iron?

Model answer: Oxygen does not directly oxidize iron in the way combustion does. Instead, dissolved oxygen serves as the cathodic reactant: O2 + 2H2O + 4e- -> 4OH-. This oxygen reduction reaction consumes the electrons produced by the anodic dissolution of iron (Fe -> Fe2+ + 2e-). Without oxygen (or another cathodic reactant like H+), the electrons produced by iron dissolution would accumulate at the metal surface, polarizing it cathodically and slowing the anodic reaction to a near halt. Dissolved oxygen depolarizes the cathode by providing an electron sink, allowing the anodic dissolution to continue at a high rate. This is why deaeration (removing dissolved oxygen from water) is a primary corrosion control strategy in boilers, cooling water systems, and oil production facilities.

This illustrates the electrochemical nature of corrosion: both an anodic reaction (metal dissolution) and a cathodic reaction (electron consumption) must occur simultaneously. The slower of the two reactions controls the overall corrosion rate. In near-neutral aerated water, the cathodic reaction (oxygen reduction) is typically rate-limiting, so the corrosion rate is proportional to the rate of oxygen transport to the metal surface. This is why stagnant aerated water often corrodes steel faster in crevices where oxygen is locally depleted — not because the crevice itself is aggressive, but because differential aeration creates a galvanic cell between the oxygen-rich surface (cathode) and the oxygen-depleted crevice (anode).