Isostructural species are defined as those having the same molecular geometry . We can determine the geometry of each species using the Valence Shell Electron Pair Repulsion (VSEPR) theory:

1. $SiF_{4}$ and $SF_{4}$: Silicon (Group 14) has four valence electrons, which it uses to form four bonds with fluorine atoms, resulting in a tetrahedral geometry . Sulphur (Group 16) has six valence electrons; it forms four bonds with fluorine and retains one lone pair ($AB_{4}E$ type), resulting in a see-saw geometry . Because their geometries differ, they are not isostructural.
2. $IO_{3}^{-}$ and $XeO_{3}$: Iodine (with a negative charge) and Xenon both effectively have eight valence electrons. They each form three bonds with oxygen and have one lone pair ($AB_{3}E$ type), leading to a trigonal pyramidal geometry .
3. $BH_{4}^{-}$ and $NH_{4}^{+}$: Both central atoms have four valence electrons involved in bonding (Boron gains one for the negative charge; Nitrogen loses one for the positive charge). Both form four bonds with hydrogen, resulting in tetrahedral geometries .
4. $PF_{6}^{-}$ and $SF_{6}$: Both phosphorus (with a negative charge) and sulphur effectively have six valence electrons, forming six bonds with fluorine atoms. This results in an octahedral geometry for both species .

Therefore, the only pair that is not isostructural is $SiF_{4}$ and $SF_{4}$.