Corrosion, Galvanization and Sacrificial Protection Explained

It is about which metal loses electrons more easily. Zinc has a more negative standard electrode potential (E° for Zn²⁺/Zn = −0.76 V) than iron (E° for Fe²⁺/Fe = −0.44 V). More negative means zinc is oxidised (corrodes) more easily. So when zinc and iron are joined, zinc becomes the anode and corrodes first, saving the iron. If you tried to coat iron on zinc, the iron would be the more easily corroded piece is FALSE — iron is less reactive, so it would not sacrifice itself, and the zinc underneath would rust away. This is exactly why the reverse coating does not work. This idea is a direct NEET 2016 question.

Sacrificial protection means we connect the metal we want to save (like iron) to a more reactive metal (like zinc or magnesium). The more reactive metal loses electrons first and corrodes, acting as a sacrificial anode. It gets used up over time, but the main object stays protected. Ships and underground pipes use blocks of magnesium or zinc for this reason. The reactive metal is the anode; the protected iron becomes the cathode, so it does not oxidise.

They are linked but not identical. Galvanization means coating iron with a layer of zinc. This works in two ways: first, the zinc layer forms a physical barrier that keeps air and water away from iron. Second, even if the zinc layer gets scratched and iron is exposed, the zinc still corrodes first because it is the anode — that part is sacrificial protection. So galvanization uses sacrificial protection as its backup mechanism. For NEET, remember galvanization = zinc coating, and its protection continues even after the coating breaks.

Rusting is an electrochemical process. At one spot on the iron (the anode) iron loses electrons: Fe → Fe²⁺ + 2e⁻. Water acts as the electrolyte that carries the charge, and at another spot (the cathode) oxygen from air takes the electrons: O₂ + 4H⁺ + 4e⁻ → 2H₂O. Without water there is no electrolyte to complete the circuit, and without oxygen the cathode reaction cannot happen. That is why iron does not rust in dry air or in pure water free of oxygen. Rust itself is hydrated ferric oxide, Fe₂O₃·xH₂O.

Tin is less reactive than iron (tin has a less negative electrode potential). So in a tin-coated iron, tin acts as a barrier only. As long as the tin layer is perfect, iron is safe. But once the tin layer is scratched and iron is exposed, iron becomes the anode (because iron is now the more reactive of the two) and rusts even faster. Zinc is the opposite: even after scratching, zinc keeps sacrificing itself. This is why galvanized (zinc) iron is safer than tinned iron for long-term protection.

Corrosion is the slow eating away of a metal when it reacts with air and moisture and turns into its oxide or salt. Rusting of iron, tarnishing of silver (black coating), and the green layer on copper are all types of corrosion. In corrosion the metal is oxidised — it loses electrons.

Rust is hydrated ferric oxide, written as Fe₂O₃·xH₂O. It forms when iron first turns into Fe²⁺ ions, which are then oxidised further by oxygen to Fe³⁺, and this combines with water to give the brown-red rust.

Zinc and magnesium are the most common. They are more reactive than iron (more negative electrode potential), so they corrode first and protect the iron object. Magnesium blocks are often bolted onto ships and buried near steel pipelines.

It greatly slows rusting but does not last forever. The zinc coating acts as a barrier and also as a sacrificial anode if scratched. Over time the zinc is used up, and after that the iron can start to rust. Still, it protects iron far longer than no coating.

NEET asks corrosion as an electrochemistry application. You must know it is an electrochemical process (anode and cathode form on the metal surface), why zinc protects iron using electrode potential values, and the meaning of sacrificial protection and galvanization. These are common one-mark reasoning questions.