What Corrosion Is
Corrosion slowly coats the surfaces of metallic objects with oxides or other salts of the metal. The rusting of iron, the tarnishing of silver, and the development of a green coating on copper and bronze are all examples of corrosion. It causes enormous economic damage to buildings, bridges, ships and to every object made of metal — especially of iron — and crores of rupees are lost every year on this account.
In essence, corrosion is a process in which a metal is oxidised by loss of electrons to oxygen, with the formation of oxides. The corrosion of iron, commonly known as rusting, occurs in the presence of both water and air. While the detailed chemistry is complex, NCERT stresses that rusting is best understood as an electrochemical process rather than a simple chemical one — and that single idea unlocks the entire topic.
| Metal | Corrosion product | Common name / appearance |
|---|---|---|
| Iron (Fe) | $\ce{Fe2O3.\mathit{x}H2O}$ | Rust — flaky, reddish-brown |
| Silver (Ag) | $\ce{Ag2S}$ | Tarnish — black coating |
| Copper (Cu) | basic copper carbonate | Patina — green coating |
Rusting as an Electrochemical Cell
The key insight is that iron does not corrode uniformly. At a particular spot on the surface — typically where the metal is strained, impure, or covered by a thicker water film — oxidation takes place, and that spot behaves as the anode. Electrons released there travel through the body of the metal to another spot, where they reduce oxygen; that spot behaves as the cathode. A thin film of moisture containing dissolved ions bridges the two spots and acts as the electrolyte.
All four ingredients of a galvanic cell are therefore present: an anode, a cathode, an electronic conductor (the metal) and an ionic conductor (the moisture film). The metal effectively short-circuits a miniature galvanic cell onto itself. This is exactly why the topic sits inside the galvanic cell framework, and why electrode potentials predict where corrosion will and will not occur.
Electrochemical model of rusting on an iron surface — separated anodic and cathodic regions bridged by a water film.
The Anode and Cathode Half-Reactions
At the anodic spot, iron passes into solution as ferrous ions, releasing electrons:
$$\ce{Anode:\quad 2Fe(s) -> 2Fe^{2+}(aq) + 4e^-} \qquad E^\circ_{(\ce{Fe^{2+}/Fe})} = -0.44\ \text{V}$$
The electrons liberated at the anode move through the metal to a cathodic spot, where they reduce dissolved oxygen. The hydrogen ions needed here are believed to come largely from carbonic acid, $\ce{H2CO3}$, formed by the dissolution of atmospheric $\ce{CO2}$ in the moisture film; dissolution of other acidic oxides also contributes $\ce{H+}$:
$$\ce{Cathode:\quad O2(g) + 4H^+(aq) + 4e^- -> 2H2O(l)} \qquad E^\circ_{(\ce{H+}\,|\,\ce{O2}\,|\,\ce{H2O})} = 1.23\ \text{V}$$
Hydrogen is not evolved at the cathode in rusting
A common error is to write the cathode reaction as $\ce{2H+ + 2e- -> H2}$. In atmospheric rusting the cathodic process is the reduction of oxygen, not of $\ce{H+}$ to hydrogen gas. The $\ce{H+}$ ions are consumed alongside $\ce{O2}$ to form water. This is precisely why corrosion needs oxygen and stops in its absence.
Cathode in rusting: $\ce{O2 + 4H+ + 4e- -> 2H2O}$ — oxygen is reduced, not hydrogen ion.
The Overall Reaction and Rust Formation
Adding the two half-reactions gives the overall cell reaction for the corrosion of iron, with a strongly positive standard cell potential that confirms its spontaneity:
$$\ce{2Fe(s) + O2(g) + 4H^+(aq) -> 2Fe^{2+}(aq) + 2H2O(l)} \qquad E^\circ_{cell} = 1.67\ \text{V}$$
The cell potential follows directly from $E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode} = 1.23 - (-0.44) = 1.67\ \text{V}$. This large positive value links straight to the thermodynamics covered under the electrochemical series: iron lies well above hydrogen and below oxygen, so it is readily oxidised in a moist, aerated environment.
The story does not end at $\ce{Fe^{2+}}$. The ferrous ions are further oxidised by atmospheric oxygen to ferric ions, which come out of solution as rust — hydrated ferric oxide, accompanied by further production of hydrogen ions:
$$\ce{2Fe^{2+}(aq) + 2H2O(l) + 1/2 O2(g) -> Fe2O3(s) + 4H^+(aq)}$$
The deposited solid is best written as $\ce{Fe2O3.\mathit{x}H2O}$, where $x$ is variable. Because rust is porous and flaky rather than adherent, it offers no barrier to the underlying metal: fresh iron is continually exposed and the corrosion eats inward. This is the structural difference from a protective oxide such as that on aluminium, which seals the surface.
Show that the standard cell potential for the rusting of iron is 1.67 V, given $E^\circ_{(\ce{Fe^{2+}/Fe})} = -0.44\ \text{V}$ and $E^\circ$ for $\ce{O2 + 4H+ + 4e- -> 2H2O}$ is $1.23\ \text{V}$.
The cathode is the reduction of oxygen ($E^\circ = 1.23\ \text{V}$) and the anode is the oxidation of iron (reduction potential $-0.44\ \text{V}$). Using $E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$:
$$E^\circ_{cell} = 1.23 - (-0.44) = 1.67\ \text{V}$$
The positive value confirms the corrosion reaction is spontaneous, with $\Delta_r G^\circ = -nFE^\circ_{cell} < 0$.
Role of Oxygen, Water and Electrolyte
Three environmental components govern whether and how fast iron rusts. Each maps onto a distinct part of the electrochemical model.
| Component | Electrochemical role | Consequence if absent |
|---|---|---|
| Water ($\ce{H2O}$) | Forms the moisture film that acts as the electrolyte; medium that supplies $\ce{H+}$ and dissolves $\ce{O2}$ | No electrolyte → no ion transport → rusting halts in dry air |
| Oxygen ($\ce{O2}$) | Species reduced at the cathode; later oxidises $\ce{Fe^{2+}}$ to $\ce{Fe^{3+}}$ | No cathodic reaction → iron does not rust in deaerated water |
| Dissolved ions (electrolyte) | Raise conductivity of the film, completing the internal circuit faster | Pure water film → high resistance → very slow corrosion |
The dependence on both water and oxygen is the experimentally observed fact NCERT highlights: iron does not rust appreciably in dry air, nor in water from which oxygen has been completely removed. It is only the simultaneous presence of moisture and air that drives the coupled anodic and cathodic processes.
Corrosion is a galvanic cell in disguise. Revise how anode, cathode and EMF are defined in Galvanic Cells & Electrode Potential before the half-reactions here.
Factors Accelerating Corrosion
Anything that strengthens the anode, the cathode or the electrolyte path speeds up corrosion. The most heavily tested factors are summarised below.
| Factor | Why it accelerates rusting |
|---|---|
| Presence of moisture | Provides the electrolyte film; humid climates rust iron faster than dry ones |
| Dissolved electrolytes (salts, acids) | Increase conductivity of the film; sea water and acid rain are highly corrosive |
| Acidic atmosphere ($\ce{CO2}$, $\ce{SO2}$ oxides) | Supply $\ce{H+}$ that drives the cathodic oxygen-reduction reaction |
| Impurities / strained metal | Create well-defined anodic and cathodic sites, localising attack |
| Contact with a less active metal | Iron becomes the anode of the couple and corrodes preferentially |
Salt water is corrosive because of conductivity, not just chemistry
Students often attribute the rapid rusting of iron near the sea to chloride "attacking" iron directly. The dominant electrochemical reason is simpler: dissolved salts make the moisture film a far better electrolyte, so the internal ionic circuit of the corrosion cell carries current much more readily and the coupled anode–cathode reactions run faster.
Prevention of Corrosion
Preventing corrosion is of prime importance: it not only saves money but also averts accidents such as a bridge collapse or the failure of a critical component. The strategies fall into two broad families — barrier methods that exclude the environment, and electrochemical methods that manipulate which metal acts as the anode.
Barrier (surface) protection
The simplest approach is to stop the metal surface from coming into contact with the atmosphere. This is done by covering the surface with paint or with certain chemicals (for example bisphenol-based anti-rust coatings). A surface coating of another metal that is inert or that itself reacts to save the object — such as tin (Sn) or zinc (Zn) — is equally common. Coating iron with zinc is specifically called galvanisation.
| Method | How it works | Behaviour when scratched |
|---|---|---|
| Painting / oil / grease | Physical barrier excluding $\ce{O2}$ and $\ce{H2O}$ | Rusting begins at the exposed gap |
| Tin plating (tinning) | Inert barrier; Sn is less easily oxidised than Fe | Iron becomes anode → rusts faster at the scratch |
| Galvanisation (zinc coating) | Barrier and sacrificial action; Zn is more easily oxidised than Fe | Zinc corrodes sacrificially → iron still protected |
Tin protects only while intact; zinc protects even when broken
Because tin sits below iron in ease of oxidation, a scratched tin can rusts rapidly — the exposed iron becomes the anode of a tin–iron couple. Zinc sits above iron, so galvanised iron remains protected even at a scratch, because zinc continues to corrode in place of the iron.
Galvanised (Zn) = sacrificial, fail-safe. Tinned (Sn) = barrier only, fail-dangerous.
Sacrificial and Cathodic Protection
The most elegant electrochemical method is to provide a sacrificial electrode of a more active metal — typically magnesium (Mg) or zinc (Zn) — connected to the iron object. The active metal, having a more negative reduction potential, becomes the anode and corrodes itself, while the iron is forced to act as the cathode and is preserved. Because the protected metal is made the cathode, this is also called cathodic protection.
This is the principle used to protect underground pipelines, ship hulls and storage tanks: blocks of magnesium or zinc are bolted on and periodically replaced as they dissolve. The iron survives as long as the sacrificial anode lasts.
Cathodic (sacrificial) protection — a magnesium block is wired to a buried iron pipe and corrodes in its place.
Note the contrast in electron direction between Figure 1 and Figure 2. In unprotected rusting, electrons flow within the iron from its own anodic to its own cathodic spots. In cathodic protection, electrons are pumped into the iron from an external sacrificial anode, holding the iron at a potential where it cannot oxidise. This same principle of using a metal's position in the activity sequence to dictate the direction of electron flow is what powers a working battery — only there it is harnessed deliberately to do useful work.
Corrosion of Other Metals
Although iron dominates the NEET syllabus, the same electrochemical logic explains the corrosion of other metals. Silver tarnishes to black silver sulphide; copper and bronze acquire a green basic-carbonate patina. The crucial difference lies in the nature of the product: the patina on copper and the oxide layer on aluminium are adherent and self-protecting, whereas rust on iron is porous and offers no protection — which is why iron, though not the most reactive metal, suffers the worst practical corrosion damage.
This contrast matters for exam answers. A protective oxide forms a continuous, adherent film that seals the surface; aluminium owes its durability to such a self-healing $\ce{Al2O3}$ layer. Iron's oxide instead cracks and flakes, exposing fresh metal, so corrosion never self-limits. That is the practical reason every prevention strategy on the syllabus targets either the supply of water and oxygen or the identity of the anode, rather than relying on rust to act as a barrier.
Corrosion in one screen
- Corrosion is an electrochemical process: iron sets up a galvanic cell on its own surface, with separated anodic and cathodic spots bridged by a moisture film.
- Anode: $\ce{2Fe -> 2Fe^{2+} + 4e^-}$ ($E^\circ = -0.44\ \text{V}$). Cathode: $\ce{O2 + 4H+ + 4e- -> 2H2O}$ ($E^\circ = 1.23\ \text{V}$). Overall $E^\circ_{cell} = 1.67\ \text{V}$.
- $\ce{Fe^{2+}}$ is further oxidised by $\ce{O2}$ to rust — hydrated ferric oxide $\ce{Fe2O3.\mathit{x}H2O}$, which is porous and non-protective.
- Rusting needs both water and oxygen; dissolved salts and acidic oxides accelerate it by raising film conductivity and supplying $\ce{H+}$.
- Prevention: barrier coatings (paint, tin), galvanisation (zinc coating), and sacrificial / cathodic protection using Mg or Zn anodes.
- Zinc protects iron even when scratched (sacrificial); tin accelerates rusting once scratched (barrier only).