Hybridisation theory has failed in a few aspects, notably in explaining the energy considerations for the involvement of $$d-orbitals$$ in chemical bonding. This can be $$well-explained$$ by means of an example. Consider, for instance, how the theory in question accounts for the bonding in phosphorus pentachloride (PCl5). The $$d-orbitals$$ are large, comparatively distant from the nucleus and high in energy. Radial distances of orbitals from the nucleus seem to reveal that $$d-orbitals$$ are far too high in energy to 'mix' with $$s-$$ and $$p-orbitals$$. $$3s$$ $$-$$ $$0.47$$ , $$3p$$ $$-$$ $$0.55$$, $$3d$$ $$-$$ $$2.4$$ (in$$ $$angstroms). Thus, at first glance, $$sp3d$$ hybridisation seems improbable. However, a closer examination of the factors that affect orbital size (and$$ $$energy) reveals more. Formal charge on the central atom is one such factor, and it is obvious that the P atom in $$PCl5$$ carries quite a large partial positive charge. Thus the $$3d$$ orbital contracts in size to such an extent that hybridisation with s and p orbitals may occur. Further, note the cases in which $$d-orbital$$ participation was proposed in hybridisation: $$SF6(sulfur$$ $$hexafluoride)$$, $$IF7$$, $$XeF6$$; in all these molecules, the central atom is surrounded by the highly electronegative fluorine atom, thus making hybridisation probable among s, p and d orbitals. A further study reveals that orbital size also depends on the number of electrons occupying it. The least stable molecule among the following $$isSH6$$ does not exist becauseThe most probable $$element(X)$$ with which silicon can form an ion of the stoichiometry $$SiX62-$$ and the hybridization of silicon in it is

Question

Hybridisation theory has failed in a few aspects, notably in explaining the energy considerations for the involvement of $$d-orbitals$$ in chemical bonding. This can be $$well-explained$$ by means of an example. Consider, for instance, how the theory in question accounts for the bonding in phosphorus pentachloride (PCl5). The $$d-orbitals$$ are large, comparatively distant from the nucleus and high in energy. Radial distances of orbitals from the nucleus seem to reveal that $$d-orbitals$$ are far too high in energy to 'mix' with $$s-$$ and $$p-orbitals$$. $$3s$$ $$-$$ $$0.47$$ , $$3p$$ $$-$$ $$0.55$$, $$3d$$ $$-$$ $$2.4$$ (in$$ $$angstroms). Thus, at first glance, $$sp3d$$ hybridisation seems improbable. However, a closer examination of the factors that affect orbital size (and$$ $$energy) reveals more. Formal charge on the central atom is one such factor, and it is obvious that the P atom in $$PCl5$$ carries quite a large partial positive charge. Thus the $$3d$$ orbital contracts in size to such an extent that hybridisation with s and p orbitals may occur. Further, note the cases in which $$d-orbital$$ participation was proposed in hybridisation: $$SF6(sulfur$$ $$hexafluoride)$$, $$IF7$$, $$XeF6$$; in all these molecules, the central atom is surrounded by the highly electronegative fluorine atom, thus making hybridisation probable among s, p and d orbitals. A further study reveals that orbital size also depends on the number of electrons occupying it. The least stable molecule among the following $$isSH6$$ does not exist becauseThe most probable $$element(X)$$ with which silicon can form an ion of the stoichiometry $$SiX62-$$ and the hybridization of silicon in it is

Infinity Learn · Accepted Answer

Bulky groups around the central atoms create repulsions.In simple words we can say that the overlap of the $$d-orbitals$$ and the $$s-orbitals$$ is poor.Vacant orbitals of silicon also participates.

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