Sulfur hexafluoride (\(\text{SF}_6\)) is a colorless, odorless gas frequently used in the electrical industry as a high-voltage insulator. This molecule is composed of a single sulfur atom and six fluorine atoms. Although \(\text{SF}_6\) contains bonds that are highly unequal in their electron sharing, it is a nonpolar molecule. This conclusion stems from a careful evaluation of the molecule’s internal charge separation and its three-dimensional shape.
Defining Molecular Polarity
To determine the overall electrical nature of any molecule, it is necessary to first distinguish between bond polarity and molecular polarity. Bond polarity describes the unequal sharing of electrons between two atoms due to a difference in their electronegativity. Electronegativity measures an atom’s ability to attract the shared electrons in a covalent bond.
The \(\text{S-F}\) bonds within sulfur hexafluoride are polar because fluorine is significantly more electronegative than sulfur. This difference causes the electron density to be pulled closer to the fluorine atoms, creating a slight negative charge (\(\delta^-\)) on each fluorine and a corresponding slight positive charge (\(\delta^+\)) on the central sulfur atom. This charge separation is referred to as a bond dipole moment.
Molecular polarity is a measurement of the charge distribution across the entire molecule, represented by the net dipole moment. A molecule is considered polar only if it possesses a net dipole moment greater than zero. This overall moment is essentially the vector sum of all the individual bond dipole moments. For a molecule to be nonpolar, the individual polar bonds must be arranged in a way that their effects perfectly neutralize each other.
The Octahedral Structure of \(\text{SF}_6\)
The shape of sulfur hexafluoride is dictated by the Valence Shell Electron Pair Repulsion (VSEPR) theory. This theory states that electron groups around a central atom will arrange themselves as far apart as possible to minimize repulsive forces.
In \(\text{SF}_6\), the central sulfur atom is bonded to six fluorine atoms and has no non-bonding lone pairs of electrons. This configuration requires the six \(\text{S-F}\) bonds to spread out as evenly as possible in space. The resulting molecular geometry is a highly symmetrical shape known as an octahedron.
This octahedral structure places the six fluorine atoms at the corners of an imaginary eight-sided figure, centered on the sulfur atom. The atoms are positioned at perfect \(90^\circ\) angles relative to their neighbors, and three pairs of fluorine atoms are directly opposite each other at \(180^\circ\). This highly regular arrangement is the direct cause of the molecule’s nonpolar nature.
How Symmetry Causes Zero Polarity
Even though each of the six individual \(\text{S-F}\) bonds is polar, the arrangement of these polar bonds cancels out their effects completely. Each bond dipole moment acts as a vector, a quantity that has both magnitude and direction.
In the \(\text{SF}_6}\) molecule, every \(\text{S-F}\) bond dipole moment is exactly opposed by an equal and opposite \(\text{S-F}\) bond dipole moment directly across the central sulfur atom. For example, the upward-pulling dipole moment of the fluorine atom at the top of the octahedron is neutralized by the downward-pulling dipole moment of the fluorine atom at the bottom. The same cancellation occurs horizontally and side-to-side.
When all these opposing vectors are added together, the total vector sum is zero. This complete cancellation of charge separation across the molecule means there is no overall positive or negative end to the structure. Therefore, the perfect octahedral symmetry of sulfur hexafluoride ensures it is a nonpolar molecule.