Given two statements: Statement I: In Cl2 molecule the covalent radius is double of the atomic radius of chlorine.…

🧠 Covalent radius is HALF the bond distance, not double the atomic radius — Statement I has the relation upside down Two different size measures get confused here. Covalent radius is defined for bonded atoms: half the internuclear distance. Atomic (Van der Waals) radius is defined for non-bonded atoms in close packing: half the closest-approach distance between non-bonded neighbours. Covalent < Van der Waals, always.

🗺️ Statement I check "In $\ce{Cl2}$ molecule, the covalent radius is double the atomic radius of chlorine." The covalent radius is half the bond length, and atomic (Van der Waals) radius is bigger than the covalent radius. So saying "covalent radius is double atomic radius" reverses the truth. False.

Statement II check "Radius of anionic species is always greater than their parent atomic radius." When you add an electron, nuclear charge stays the same but electron-electron repulsion increases. The cloud expands. So $\ce{Cl-}$ ($\approx 181\,pm$) is larger than $\ce{Cl}$ atom ($\approx 99\,pm$ covalent radius). True for every parent → anion conversion. True.

⚠️ The trap Students often misremember the covalent vs Van der Waals relation. Memorize this one line: "Covalent radius is for bonded atoms (smaller). Van der Waals is for non-bonded atoms (bigger)." Anything that swaps the two is wrong. The factor "double" in Statement I is the giveaway — covalent radius equals half the bond distance, never double anything.

$\boxed{\text{Answer: (c)}}$