When atoms form a molecule, the sharing of electrons determines many physical properties, such as solubility and reactivity. This sharing is rarely equal, leading to molecular polarity. Polarity is quantified by the dipole moment, a vector quantity reflecting the overall distribution of positive and negative charge centers. For sulfur tetrafluoride (\(\text{SF}_4\)), the chemical formula does not reveal if it is polar or nonpolar. The answer depends entirely on the balance between the polarity of its individual bonds and its precise three-dimensional shape, which dictates whether internal electrical forces cancel out.
Understanding Molecular Polarity
Molecular polarity arises from two factors: the polarity of individual bonds and the molecule’s geometry. Bond polarity is based on electronegativity, an atom’s power to attract electrons within a chemical bond. In the \(\text{S-F}\) bond, Fluorine (F) has a significantly higher electronegativity (3.98) than Sulfur (S) (about 2.58).
This difference means shared electrons are pulled closer to the Fluorine atom, creating a polar covalent bond. The Fluorine end acquires a partial negative charge (\(\delta^-\)), and the Sulfur end acquires a partial positive charge (\(\delta^+\)), forming a bond dipole. A molecule’s overall polarity, or net dipole moment, is the vector sum of all these individual bond dipoles. For a molecule to be nonpolar, these vectors must perfectly cancel out, which occurs only in highly symmetrical shapes, such as methane (\(\text{CH}_4\)) or carbon dioxide (\(\text{CO}_2\)). If the molecular structure is asymmetrical, the dipoles will not cancel, resulting in a net dipole moment and a polar molecule.
Determining the Molecular Geometry of \(\text{SF}_4\)
The polarity of \(\text{SF}_4\) depends entirely on its precise three-dimensional structure. To determine this, we count the valence electrons: Sulfur contributes six, and the four fluorine atoms contribute seven each, totaling 34 valence electrons.
Four electrons form the four single bonds connecting sulfur to the fluorine atoms. After accounting for the three lone pairs on each fluorine atom, a single lone pair remains on the central sulfur atom.
The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the arrangement of these electron regions. The sulfur atom has five electron density regions: four bonding pairs and one lone pair, classifying it as an \(\text{AX}_4\text{E}_1\) structure. These five regions arrange themselves into a trigonal bipyramidal electron geometry to maximize distance.
The molecular geometry—the shape defined only by the atoms—is not trigonal bipyramidal because the lone pair occupies space. Lone pairs exert a greater repulsive force than bonding pairs, forcing the lone pair into the less crowded equatorial position. This placement results in an asymmetrical molecular shape known as the “seesaw” or disphenoidal geometry.
Why \(\text{SF}_4\) Possesses a Dipole Moment
The asymmetrical “seesaw” molecular geometry is why \(\text{SF}_4\) has a net dipole moment. The four \(\text{S-F}\) bonds are polar, with electron density pulled toward the fluorine atoms. If the molecule were perfectly symmetrical, like octahedral \(\text{SF}_6\), the bond dipoles would cancel out, resulting in a nonpolar molecule.
In the seesaw shape, the axial and equatorial fluorine atoms are arranged unevenly. The bond dipoles do not perfectly oppose each other. For example, the axial \(\text{F-S-F}\) angle is compressed to about \(173^\circ\) (instead of \(180^\circ\)), and the equatorial \(\text{F-S-F}\) angle is reduced to about \(102^\circ\) (instead of \(120^\circ\)) due to the lone pair’s repulsion.
This structural distortion prevents the individual bond dipoles from canceling completely. The lone pair on the sulfur atom also contributes a significant dipole moment, pointing away from the central atom. The vector addition of these non-canceling dipoles results in a net, non-zero dipole moment. This net charge separation makes \(\text{SF}_4\) a polar molecule, with an experimentally measured dipole moment of approximately 0.632 Debye.