Q.Why do noble gases have comparatively large atomic sizes?

Step 1: How atomic radius is normally measured.

For most elements, atomic radius is measured as a covalent radius — half the distance between the nuclei of two atoms joined by a single covalent bond.

Step 2: Noble gases don't form such bonds.

Since noble gas atoms (under ordinary conditions) do not form covalent bonds with each other (they exist as free, monatomic species), there is no covalent-bond distance to measure. Instead, their atomic radius is measured as a van der Waals radius — half the distance between the nuclei of two adjacent, non-bonded (touching) atoms in the solid state, held together only by weak van der Waals (dispersion) forces.

Step 3: Why van der Waals radius is larger.

In a covalent bond, the bonding electron pair is drawn into the region between the two nuclei, pulling the nuclei relatively close together. In a van der Waals contact, there is no such shared bonding electron density pulling the atoms together — the atoms simply touch at the outer edge of their electron clouds, held apart by inter-electronic repulsion, resulting in a much larger internuclear distance.

Step 4: Conclusion. …

Explain that atomic radius comparisons across the periodic table normally use covalent radii, but noble gases (forming no bonds) can only be measured as van der Waals radii, which are inherentl …

• Do not attribute the large size to noble gases having 'more electrons' or being 'at the end of the period' without the specific van der Waals-vs-covalent radius distinction, which is the actual textbook reasoning.
• Do not imply noble gas atoms are literally larger in an absolute/electronic sense than the halogen before them — the apparent jump in size is a consequence of comparing two different TYPES of radius, not a real discontinuity in atomic structure. …