Question bank — Fajan's rules — covalent character in ionic compounds

False. Any cation exerts some field on any anion, so there is always non-zero polarisation; "ionic" is one end of a continuous scale, never a pure state.

True. The polarising field goes as $1/r^2$ from Coulomb's law and electric fields, so shrinking the cation intensifies the pull on the anion cloud.

False. Opposite role — a larger anion holds outer electrons farther from its nucleus, making them loosely held and more polarisable, so bigger anion means more covalent.

False. Charge enters the field linearly ($E\propto q$) but radius enters as $1/r^2$, so halving the radius (roughly $\times 4$) outweighs doubling the charge ($\times 2$).

False. It is only a rough proxy; the real field $\propto q/(r_{\text{cation}}+r_{\text{anion}}{)}^2$ uses the full separation and a square, and $\varphi$ ignores the anion entirely.

True. The extra electron on a $2-$ anion is held more loosely per unit nuclear charge, adding to the size effect, so sulphides tend to be more covalent than chlorides.

False. Rule 4 breaks the tie — Cu⁺'s $3d^{10}$ core shields poorly, so its effective pull on the anion exceeds Na⁺'s, making CuCl far more covalent than NaCl.

True. Covalency replaces the strong infinite ionic lattice with more discrete molecular units, so less energy is needed to melt — hence AlCl₃ sublimes while NaCl melts near 801 °C.

False. The colour arises from Charge-transfer transitions and colour / band-structure effects; polarisation only correlates with the trend, it is not the mechanism.

True. Covalent character weakens ion–dipole hydration, so the balance in Solubility and lattice/hydration energy shifts toward the solid — sulphides and AgI are classic poorly-water-soluble examples.

Wrong for the cation — bigger cation means a weaker field ($1/r^2$), hence less covalent. "Bigger = more covalent" applies only to the anion.

Charge matters too — a $2-$ anion holds its electrons more loosely than a comparable $1-$ anion, so O²⁻/S²⁻ are more polarisable than F⁻/Cl⁻.

The true field uses $q/(r_{\text{cation}}+r_{\text{anion}}{)}^2$; the squared, separation-inclusive denominator can dominate, so size is not automatically subordinate to charge.

$\varphi =q_{\text{cation}}/r_{\text{cation}}$ contains no anion term at all — it misses anion size, anion charge (Rule 3) and cation configuration (Rule 4).

True for those two, but comparing them to Zn²⁺/Cd²⁺/Hg²⁺ (pseudo-noble-gas $d^{10}$) requires Rule 4 — the $d^{10}$ ions polarise more at equal charge and size.

Electronegativity of the element is not the driver here; I⁻ is huge and squishy while F⁻ is tiny and tight, so AlI₃ is the more covalent one.

The relevant comparison is lattice vs discrete molecules — molecular covalent solids have weak intermolecular forces, so covalent character lowers melting/boiling points.

Polarisability is asymmetric — the big, loosely-bound anion cloud distorts; the small, tightly-held cation is the distorter, not the distorted.

Because the cation acts like a point charge and Coulomb's law and electric fields gives $E=kq/r^2$; the inverse-square makes small changes in separation matter enormously.

Diffuse, non-spherical $d$ electrons shield the nucleus poorly, so the anion feels a larger effective nuclear charge than the ion's formal $+1$ or $+2$ suggests.

S²⁻ is large and $2-$, so it is highly polarisable; strong polarisation gives covalent character that both darkens the solid (charge-transfer) and lowers hydration-driven solubility.

It compresses "high charge, small cation" into one comparable number for quick ranking — useful as a first pass, provided you remember it omits the anion and the $1/r^2$ dependence.

The distortion drags anion electron density into the region between the two nuclei, and shared inter-nuclear density is exactly what a covalent bond is.

The visible absorption comes from charge-transfer / band-structure transitions in the solid; polarisation trends run parallel to it but do not by themselves set the transition energy.

Minimal polarisation — weak field (large $r$, only $+1$) plus a tight little anion — so CsF is about as close to ideal ionic (Ionic bonding — lattice energy) as a real salt gets.

The one with the larger and/or higher-charged anion, because equal $\varphi$ says nothing about the anion, and Rule 3 then decides via polarisability.

No — polarisation is a mutual geometric effect; with nothing to distort and no separation to define, the concept simply does not apply.

MgCl₂ — Mg²⁺ is the smaller cation, so its field is stronger, giving greater polarisation and covalent character than the larger Ca²⁺.

A (nearly) pure covalent bond — the "ionic with covalent character" description has slid all the way to the covalent end of the continuum, i.e. maximum electron sharing.