Basic idea of rusting

Rusting of iron is an electrochemical process. Tiny anodic and cathodic regions develop on the iron surface in presence of moisture and oxygen.

• Anode: $\mathrm{Fe \rightarrow Fe^{2+} + 2e^-}$
• Cathode (in neutral/alkaline medium): $\mathrm{O_2 + 2H_2O + 4e^- \rightarrow 4OH^-}$

Then: $\mathrm{Fe^{2+} + 2OH^- \rightarrow Fe(OH)_2}$ which further gets oxidized to hydrated ferric oxide, i.e. rust.

Check each option

Option A

"Rusting of iron is envisaged as setting up of electrochemical cell on the surface of iron object."

This is correct because rusting indeed occurs via formation of small electrochemical cells on the iron surface.

Option B

"Dissolved acidic oxides $\mathrm{SO_2, NO_2}$ in water act as catalyst in the process of rusting."

These acidic oxides dissolve in water and make the medium acidic, increasing conductivity and accelerating corrosion/rusting.

Though the wording "catalyst" is not the most precise, in standard exam context this statement is taken as correct since these dissolved gases promote rusting.

Option C

"Coating of iron surface by tin prevents rusting, even if the tin coating is peeling off."

This is not correct.

Reason: Tin is less reactive than iron. If the tin coating is damaged, then iron gets exposed and a galvanic cell forms where iron acts as the anode and corrodes faster.

So tin coating protects iron only as long as the coating remains intact. If it peels off or gets scratched, rusting is actually accelerated.

Hence, C is incorrect.

Option D

"When $\mathrm{pH}$ lies above 9 or 10, rusting of iron does not take place."

This is correct in the usual electrochemistry context. In strongly alkaline medium, rusting is greatly suppressed because iron becomes passive.

Conclusion

The statement which is not correct is: $\boxed{\text{C}}$