Statements on Covalency and Oxidation State of Oxygen

To evaluate the given statements, let us analyze them individually based on the chemical properties of oxygen and periodic trends:

Analysis of Statement I:

1. Covalency of Oxygen:

   • Oxygen belongs to the second period of the periodic table and has the valence shell electronic configuration of $2s^2 2p^4$.
   • Because it lacks vacant $d$-orbitals, its maximum coordination number or valence shell capacity is restricted to 8 electrons (utilizing one $2s$ orbital and three $2p$ orbitals). Hence, its maximum covalency is $4$, though in practice, it usually exhibits a covalency of $2$ (or $3$, as in $\text{H}_3\text{O}^+$). Thus, its covalency generally is $2$ and can go up to a maximum of $4$.

2. Oxidation State of Oxygen:

   • In $\text{SO}_2$: Oxygen is more electronegative than sulfur ($\chi_{\text{O}} > \chi_{\text{S}}$). Sulfur is in the $+4$ oxidation state, giving oxygen its typical oxidation state of $-2$.
   • In $\text{OF}_2$: Fluorine is the most electronegative element ($\chi_{\text{F}} > \chi_{\text{O}}$) and carries an oxidation state of $-1$. Consequently, for the neutral molecule $\text{OF}_2$: $\text{O.S. of O} + 2(-1) = 0 \implies \text{O.S. of O} = +2$

Therefore, Statement I is true.

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Analysis of Statement II:

• Oxygen exhibits anomalous behaviour compared to other elements of Group 16 ($\text{S}$, $\text{Se}$, $\text{Te}$, $\text{Po}$).
• This anomalous nature is primarily attributed to:
  1. Small atomic/ionic size.
  2. High electronegativity.
  3. High ionization enthalpy.
  4. Non-availability of vacant $d$-orbitals in its valence shell.

Therefore, Statement II is true.

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Conclusion:

Since both Statement I and Statement II are true, the correct option is A.