Q.Oxygen forms only OF$_2$ with fluorine while sulfur forms SF$_6$. Explain. Why ?

As with other p-block groups, the very lightest (second-period) member of groups 16 and 17 -- oxygen and fluorine respectively -- differs sharply from the rest of its own group, because of its unusually small size, unusually high electronegativity, and complete lack of accessible valence-shell d-orbitals. Oxygen is diatomic (O2) rather than polyatomic like sulfur's S8; it is paramagnetic (from two unpaired pi* electrons) while heavier chalcogens are diamagnetic; it shows only -2/-1/+2 oxidation states rather than reaching +4/+6; its hydride (water) is uniquely a liquid, thanks to strong hydrogen bonding; and its covalency is capped at about 2 (rarely 4), versus the heavier elements' ability to exceed four. Fluorine similarly diverges from the rest of group 17: its ionisation enthalpy, electronegativity and electrode potential are all higher than the group trend predicts; its radii, melting/boiling points and electron gain enthalpy are all lower than expected; its reactions are strongly exothermic; it forms only one oxoacid (HOF) instead of a whole family; and its hydride, HF, is uniquely a liquid because of strong hydrogen bonding. Both anomalies trace back to the same underlying cause: a second-period atom is small, highly electronegative, and structurally has no d-orbitals to expand into, unlike every heavier member of its own group.