Given below are two statements : Statement (I) : The oxidation state of an element in a particular compound is the charge acquired by its atom on the basis of electron gain enthalpy consideration from other atoms in the molecule. Statement (II) : $$p\pi - p\pi$$ bond formation is more prevalent in second period elements over other periods. In the light of the above statements, choose the most appropriate answer from the options given below :

We need to evaluate two statements about oxidation states and $$p\pi - p\pi$$ bonding.

Analysis of Statement I:

"The oxidation state of an element in a particular compound is the charge acquired by its atom on the basis of electron gain enthalpy consideration from other atoms in the molecule."

This statement is incorrect. The oxidation state of an element is determined based on electronegativity, not electron gain enthalpy. In the concept of oxidation state, we assume that shared electron pairs are assigned entirely to the more electronegative atom. While electron gain enthalpy and electronegativity are related concepts, they are not the same:

- Electronegativity is the tendency of an atom to attract shared electrons in a chemical bond.

- Electron gain enthalpy is the energy change when a gaseous atom gains an electron to form an anion.

Oxidation states are assigned using electronegativity rules (e.g., in H$$_2$$O, oxygen is more electronegative, so both shared pairs are assigned to O, giving it an oxidation state of $$-2$$). Using electron gain enthalpy would not always give correct results because the two properties do not always follow the same trend (for example, fluorine has a lower electron gain enthalpy than chlorine, but fluorine is more electronegative).

Analysis of Statement II:

"$$p\pi - p\pi$$ bond formation is more prevalent in second period elements over other periods."

This statement is correct. Second period elements (such as C, N, O) have small atomic sizes, which allows their 2p orbitals to overlap laterally (sideways) very effectively to form strong $$\pi$$ bonds. For elements in higher periods (third period and beyond), the p orbitals are larger and more diffuse, leading to poor lateral overlap. This is why:

- Carbon readily forms $$C=C$$ and $$C\equiv C$$ bonds (with $$p\pi - p\pi$$ bonding)

- Nitrogen forms $$N=N$$ and $$N\equiv N$$ bonds

- Oxygen forms $$O=O$$ bonds

- But silicon prefers single bonds (Si-Si) rather than double bonds, and phosphorus and sulfur rarely form $$p\pi - p\pi$$ bonds

Elements in higher periods instead use $$d\pi - p\pi$$ bonding when they form multiple bonds.

Conclusion: Statement I is incorrect, and Statement II is correct.

The correct answer is Option (3): Statement I is incorrect but Statement II is correct.