Question

Statement I: The covalency of oxygen is generally two but it can exceed up to four. The oxidation state of oxygen in SO₂ is -2 and in OF₂ it is +2.
Statement II: The anomalous behaviour of oxygen when compared to the other elements of group 16 is due to its small size and high electronegativity.
In the light of the above statements, choose the correct answer:

• A. Both Statement I and Statement II are true
• B. Both Statement I and Statement II are false
• C. Statement I is true but Statement II is false
• D. Statement I is false but Statement II is true

Answer 1

The Correct Option is D

Step 1: Understanding the Concept:
Covalency refers to the number of electron pairs shared by an atom. Oxygen belongs to the 2nd period and lacks d-orbitals, which limits its maximum covalency. Oxidation states are determined by electronegativity differences.

Step 2: Key Formula or Approach:
1. Check oxygen's covalency limits (Valence shell is $n=2$, orbitals available: $2s, 2p$). 2. Compare electronegativities of O, S, and F.

Step 3: Detailed Explanation:
1. Statement I analysis: Oxygen has only four valence orbitals ($2s$ and three $2p$). Due to the absence of d-orbitals, its covalency rarely exceeds 2 and cannot exceed 4 (it is usually limited to 3 in hydronium ions, $H_3O^+$). However, Statement I says it "can exceed up to four," which is incorrect. Oxidation states are correct (O is -2 in $SO_2$ and +2 in $OF_2$ because F is more electronegative). Since part of the statement is wrong, Statement I is false. 2. Statement II analysis: Oxygen shows anomalous behavior (like being a gas $O_2$ while others are solids, and forming H-bonds) strictly due to its high electronegativity, small size, and absence of d-orbitals. This is true.

Step 4: Final Answer:
Statement I is false, but Statement II is true.