Q.Consider the change in oxidation state of bromine corresponding to different emf values as shown in the diagram below: $\text{BrO}_4^{-} \xrightarrow{1.82\text{V}} \text{BrO}_3^{-} \xrightarrow{1.5\text{V}} \text{HBrO} \xrightarrow{1.595\text{V}} \text{Br}_2 \xrightarrow{1.0652\text{V}} \text{Br}^{-}$ Then the species undergoing disproportionation is

Corrosion — the rusting of iron, or the green film on copper and brass — is a redox process in which a metal is oxidised by atmospheric oxygen in the presence of moisture. Rusting is specifically an electrochemical process requiring a spatially separated anode and cathode on the SAME piece of metal: a water droplet on iron creates a low-oxygen anodic region beneath it, where iron dissolves ($\text{Fe} \rightarrow \text{Fe}^{2+}+2e^-$, $E^{o}=0.44$ V), and a high-oxygen cathodic region around it, where oxygen is reduced ($\text{O}_2+4\text{H}^++4e^- \rightarrow 2\text{H}_2\text{O}$, $E^{o}=1.23$ V); the overall reaction has $E^{o}=1.67$ V, confirming rusting is spontaneous. Fe²⁺ is further oxidised to Fe³⁺, which precipitates as hydrated rust, Fe2O3·xH2O.

Five protection strategies follow directly from blocking one ingredient corrosion needs: painting (blocks oxygen/moisture physically), galvanising (a sacrificial zinc coating corrodes instead of the iron beneath, since zinc is more easily oxidised), cathodic protection (a sacrificial anode of Mg or Zn, connected but not necessarily coating the whole surface, forces the protected metal to act as cathode), passivation (a strong oxidiser like conc. HNO3 forms a tough protective oxide layer, as aluminium's natural Al2O3 layer does), and alloy formation (stainless steel, an Fe–Cr alloy, resists corrosion far better than plain iron). …