NEET Practice Questions (MCQs) with Answers & Solutions

Corrosion involves the surface of metals being coated with oxides or other salts due to oxidation. Rusting of iron, tarnishing of silver, and the green coating on copper (patina) are all examples of corrosion. Electrorefining of copper is an industrial process used to purify copper, where impure copper is oxidized at the anode and pure copper is deposited at the cathode; it is not a degradation process like corrosion. (Refer to the section 'Corrosion' and 'Electrorefining' in the context).

The context states, 'Electrons released at anodic spot move through the metal and go to another spot on the metal and reduce oxygen in the presence of H+...'. Therefore, H+ ions are essential for the cathodic reaction where oxygen is reduced. Oxidation of iron occurs at the anodic spot, and oxygen is reduced at the cathodic spot. Rust ($\text{Fe}_2\text{O}_3\cdot x\text{H}_2\text{O}$) is formed by the further oxidation of $\text{Fe}^{2+}$ to $\text{Fe}^{3+}$ by atmospheric oxygen. (Refer to the detailed explanation of iron corrosion chemistry).

The text provides the standard electrode potential for the anode reaction ($\text{Fe}^{2+}/\text{Fe}$) as $E^o = -0.44 V$ and for the cathode reaction ($ \text{O}_2|\text{H}_2\text{O} \text{, } \text{H}^+$) as $E^o = 1.23 V$. The overall reaction is $\text{2Fe(s)} + \text{O}_2\text{(g)} + \text{4H}^+\text{(aq)} \longrightarrow \text{2Fe}^{2+}\text{(aq)} + \text{2H}_2\text{O(l)}$. The standard cell potential is $E^o_{cell} = E^o_{cathode} - E^o_{anode} = 1.23 V - (-0.44 V) = 1.67 V$. (Refer to the 'Corrosion' section and the equations for $E^o_{cell}$).

The text states: 'Electrons released at anodic spot move through the metal and go to another spot on the metal and reduce oxygen in the presence of H+... Cathode: O2(g) + 4 H+(aq) + 4 e– ¾® 2 H2O (l)'. This clearly indicates that $\text{H}^+$ ions are consumed to facilitate the reduction of oxygen at the cathode. The context also mentions 'further production of hydrogen ions' when ferrous ions are oxidized to ferric ions, which means they are not simply 'produced' but rather involved in a cycle. (Refer to the detailed explanation of iron corrosion chemistry).

The text explicitly states: 'The ferrous ions are further oxidised by atmospheric oxygen to ferric ions which come out as rust in the form of hydrated ferric oxide ($\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}$)'.

For a metal to act as a sacrificial electrode, it must be more easily oxidized than iron. This means it should have a more negative (or less positive) standard reduction potential than iron. From the appendix and context, iron has a standard reduction potential for $\text{Fe}^{2+}/\text{Fe}$ = -0.44 V. Comparing the options: Cu (+0.34 V), Ni (-0.23 V), Zn (-0.76 V), Ag (+0.80 V). Zinc (-0.76 V) has the most negative reduction potential among the choices, indicating it is most easily oxidized, and thus would be the most effective sacrificial electrode. (Refer to the 'Prevention of corrosion' section and Appendix III for standard potentials).

The context states: 'H+ (which is believed to be available from H2CO3 formed due to dissolution of carbon dioxide from air into water.' This indicates that dissolved carbon dioxide contributes indirectly by forming carbonic acid, which then provides the necessary $\text{H}^+$ ions for the cathodic reaction. (Refer to the explanation of the cathodic reaction in iron corrosion).

The text explicitly states: 'In corrosion, a metal is oxidised by loss of electrons to oxygen and formation of oxides.' (Refer to the 'Corrosion' section).

The text lists: 'covering the surface with paint or by some chemicals', 'cover the surface by other metals (Sn, Zn, etc.) that are inert or react to save the object', and 'provide a sacrificial electrode of another metal (like Mg, Zn, etc.) which corrodes itself but saves the object'. While applying an external electric potential might be part of cathodic protection, the specific mechanism of 'applying an external electric potential to make the object cathodic' is not explicitly mentioned as a standalone method in the way the other options are, making it the least directly supported statement from the context. (Refer to the 'Prevention of corrosion' section).

The context states: 'At a particular spot... oxidation takes place and that spot behaves as anode and we can write the reaction Anode: 2 Fe (s) ¾® 2 Fe 2+ + 4 e –'. (Refer to the 'Corrosion' section and the anode reaction equation).