Atomic and ionic radii

Periodic trends: across a period reactivity is high in group 1, lower towards the middle and highest in group 17. Down the alkali metals reactivity rises; down the halogens it falls. Trends are explained through nuclear charge and energy levels.

An atom is tiny (about 1.2 Å = 1.2 × 10⁻¹⁰ m) and its electron cloud has no sharp edge, so its radius is estimated from distances between bonded atoms.

Covalent radius (non-metals): half the single-bond length, e.g. Cl–Cl in Cl₂ is 198 pm, so r(Cl) = 99 pm. Metallic radius: half the distance between neighbouring metal cores, e.g. Cu–Cu 256 pm, so r(Cu) = 128 pm. NCERT calls either one the atomic radius.

Across a period the radius falls: electrons enter the same shell while effective nuclear charge rises and pulls them in. Period 2 (pm): Li 152, Be 111, B 88, C 77, N 74, O 66, F 64. Period 3: Na 186, Mg 160, Al 143, Si 117, P 110, S 104, Cl 99.

Down a group the radius grows: n increases, and filled inner shells shield the outer electrons. Group 1 (pm): Li 152, Na 186, K 231, Rb 244, Cs 262. Group 17: F 64, Cl 99, Br 114, I 133, At 140.

Noble gases are left out: they are monoatomic, so only non-bonded (van der Waals) radii exist, and these are much larger than covalent radii.

A cation is smaller than its atom (fewer electrons, same nuclear charge); an anion is larger (extra repulsion, lower effective nuclear charge). F 64 pm but F⁻ 136 pm; Na 186 pm but Na⁺ 95 pm.

Isoelectronic species have the same number of electrons, e.g. O²⁻, F⁻, Na⁺ and Mg²⁺ (10 each). The more protons, the smaller the ion: O²⁻ > F⁻ > Na⁺ > Mg²⁺.

Among Mg, Mg²⁺, Al and Al³⁺ the largest is Mg and the smallest is Al³⁺.