Question

The shape of \[S{{F}_{4}}\]and \[\text{Xe}{{F}_{2}}\]respectively are:
A. trigonal bipyramidal and trigonal bipyramidal
B. See-saw and linear shape
C.T-shape and linear
D.Square planar and triangular bipyramidal

Answer 1

Hint:
The first one has a molecular geometry in which the central atom is bonded to four other atoms and one lone pair. The general formula of this geometry is \[\text{A}{{\text{X}}_{4}}\text{E}\]. Here, X denotes to the four atoms bonded to the central atom and E is the lone pair.
The second one has a molecular geometry where the central atom is bonded to two other atoms and the angle between the two bonded atoms with the central one is \[{{180}^{\circ }}\].

Complete step by step answer:
The shape of \[S{{F}_{4}}\]and \[\text{Xe}{{F}_{2}}\]respectively are see-saw and linear shape.
\[S{{F}_{4}}\]:-
Valency of sulphur = 6
Valency of fluorine = 7
So, the total valence electron contribution = \[[6+(7\times 4)]\]= 34 electrons
The nearest multiple of noble gas is 32.
So, (34-32) = 2 electrons exist as a lone pair on the sulphur atom.
Now, the sulphur forms four single bonds to the fluorine atoms. Each fluorine atom has three lone pairs to complete their octet.
Since a lone pair experiences repulsion, it is placed in equatorial position to reduce repulsion. Sulphur forms four single bonds and one lone pair. So, the hybridisation is \[\text{s}{{\text{p}}^{3}}\text{d}\].
See-saw structure:


\[\text{Xe}{{F}_{2}}\]:-

The two unpaired electrons bond with the one unpaired electron of the two fluorine atoms.
So, there are five electron pairs. The two bond pairs are axial and three lone pairs are equatorial. This is because a lone pair-bond pair has more repulsion than a bond pair-bond pair.
Linear structure:

So, the correction option is B.

Note:
The structures of any compound depend on the hybridisation and the number of lone pairs and bond pairs.
Order of repulsion which determines the position of lone pairs around the central metal (either equatorial or axial):
Lone pair-lone pair > Lone pair-bond pair > Bond pair-bond pair