Q.Substantiate Lithium fluoride has the lowest solubility among group one metal fluorides.

Step 1. Solubility depends on the balance between an ionic solid's lattice enthalpy (which opposes dissolution) and the hydration enthalpy of its ions (which favours dissolution).

Step 2. Li⁺ is the smallest alkali metal cation and so has the largest-magnitude hydration enthalpy among the group 1 cations -- this is normally why lithium salts are unusually soluble (e.g. LiClO₄ up to 12× more soluble than NaClO₄).

Step 3. Fluoride, F⁻, is however also the smallest of the halide ions. Pairing the smallest cation (Li⁺) with the smallest anion (F⁻) gives an unusually SHORT, tightly-packed ionic lattice, and lattice enthalpy grows sharply as ionic separation shrinks -- so LiF has an exceptionally HIGH lattice enthalpy, larger than that of LiCl, LiBr or LiI (whose larger halide ions give a longer, less tightly bound lattice).

Step 4. Even Li⁺'s strong hydration enthalpy is not large enough to overcome this unusually high lattice enthalpy, so LiF ends up with net-unfavourable dissolution energetics -- making it poorly soluble, unlike LiCl/LiBr/LiI where the much smaller lattice enthalpy is easily outweighed by Li⁺'s hydration. …

Analyse LiF's solubility via the same lattice-enthalpy-vs-hydration-enthalpy balance used throughout the unit, noting that pairing the smallest cation …

• Explaining LiF's low solubility purely by Li+'s hydration behaviour (as if it applied to all lithium halides) without recognising that it is specifically F−'s small si …