Question

Likely bond angles of $S{F_4}$ molecules are
(A) ${89^ \circ },{117^ \circ }$
(B) ${120^ \circ },{180^ \circ }$
(C) ${45^ \circ },{118^ \circ }$
(D) ${117^ \circ },{92^ \circ }$

Answer 1

$S{F_4}$(sulphur tetrafluoride ) has trigonal bipyramidal electron geometry and see-saw as molecular geometry. The geometry is there in accordance with the valence shell electron pair repulsion theory rules. The bond angle of the given molecule depends on the position of lone pairs at different positions and the repulsions between the bond pairs and the lone pairs.

Complete answer:
We know that the sulphur tetrafluoride has the trigonal bipyramidal structure and see-saw geometry according to the valence shell electron pair repulsion theory. It is having $s{p^3}d$ hybridisation with one lone pair. The lone pair of electrons at equivalent $F-S-F$ bond angle position is less than $120$ degree. In the planar structure of sulphur tetrafluoride, the angles are less due to the bulky nature of the lone pair. The equatorial bond angles are less than $120$ due to the repulsion from the lone pairs. The axial bond angles are less than $180$ degrees due to the repulsion from the lone pairs. So likely the bond angles of $S{F_4}$ molecules are $89$ and $117$ degrees.

**Hence the correct answer is option A.

Additional information: **
Sulphur is the least electronegative element in this structure so it gets transferred in the middle of the structure and this gives us a three-dimensional view and structural information.

Note:
$S{F_4}$ molecules are polar molecules. They have five regions of electron geometry around the central sulphur atom. Here the F-S-F axial bond angles at axial point should be exactly one hundred eighty degrees due to even distribution in space but it is less than one hundred eighty degrees due to the more repulsions from the lone pairs and less from the bond pairs.