Phosphorus pentafluoride (\(\text{PF}_5\)) is a chemical compound frequently examined when learning about molecular structure and electron behavior. Determining whether a molecule is polar or nonpolar depends entirely on the molecule’s overall electronic distribution. The definitive answer is that \(\text{PF}_5\) is a nonpolar molecule, despite containing bonds that are themselves polar. This distinction highlights how a molecule’s three-dimensional shape ultimately dictates its macroscopic properties. Understanding this characteristic requires a detailed look at the forces acting within the molecule.
Defining Molecular Polarity
Molecular polarity describes the asymmetrical distribution of electric charge across a molecule, which results in a net dipole moment. This characteristic begins at the level of the individual chemical bond, where atoms share electrons unequally. The measure of an atom’s ability to attract a shared pair of electrons toward itself is called electronegativity.
A significant difference in electronegativity between two bonded atoms creates a bond dipole moment, resulting in a polar bond. In the case of phosphorus (\(\text{P}\), 2.19) and fluorine (\(\text{F}\), 3.98), the substantial difference of 1.79 indicates that the electrons are pulled strongly toward the fluorine atoms.
Consequently, each fluorine atom acquires a partial negative charge (\(\delta-\)), while the central phosphorus atom carries a partial positive charge (\(\delta+\)). Therefore, each individual \(\text{P}-\text{F}\) bond is a highly polar covalent bond. A molecule’s overall polarity, however, is the vector sum of all individual bond dipole moments.
If the molecular structure is asymmetrical, the bond dipoles will not cancel, leading to a net dipole moment and a polar molecule. Conversely, if the geometry is perfectly symmetrical, opposing bond dipoles neutralize one another, resulting in a zero net dipole moment and classifying the molecule as nonpolar overall.
Determining the Molecular Geometry of \(\text{PF}_5\)
The three-dimensional arrangement of atoms in phosphorus pentafluoride is the primary factor determining its nonpolar nature. To predict this shape, chemists use the Valence Shell Electron Pair Repulsion (VSEPR) theory. This theory posits that electron domains (bonding pairs or lone pairs) will arrange themselves around a central atom to maximize the distance between them, thus minimizing electronic repulsion.
The first step in applying VSEPR theory is determining the number of valence electrons contributed by each atom. Phosphorus (Group 15) contributes five valence electrons, and the five fluorine atoms (Group 17) contribute seven each. This totals forty valence electrons for the entire \(\text{PF}_5\) molecule.
Phosphorus acts as the central atom, forming a single bond with each of the five surrounding fluorine atoms. This arrangement utilizes all five valence electrons of the central phosphorus atom for bonding. Crucially, the central phosphorus atom has no non-bonding or lone pairs of electrons.
The central atom is surrounded by five bonding domains and zero lone pairs, designated by the \(\text{AX}_5\text{E}_0\) VSEPR notation. To accommodate five electron domains with maximum separation, the electron domain geometry and the molecular geometry are both trigonal bipyramidal.
This shape features a central phosphorus atom with five fluorine atoms positioned around it. Three fluorine atoms lie in a single plane, forming an equilateral triangle; these are the equatorial positions. The remaining two fluorine atoms occupy the axial positions, located above and below this plane along a perpendicular axis. The atoms in the equatorial plane are separated by \(120^\circ\) bond angles, while the axial atoms form \(90^\circ\) angles with the equatorial plane.
Why Symmetry Makes \(\text{PF}_5\) Nonpolar
The nonpolar character of \(\text{PF}_5\) rests entirely on the perfect symmetry of its trigonal bipyramidal structure. This precise arrangement geometrically nullifies the highly polar nature of the individual \(\text{P}-\text{F}\) bonds. Dipole moments are vector quantities, meaning they have both magnitude and direction, and must be added together using vector mathematics.
The three equatorial \(\text{P}-\text{F}\) bond dipoles are arranged in a plane at \(120^\circ\) angles. Because they are of equal magnitude and point symmetrically away from the center, their vector sum is zero. This set of three dipoles cancels itself out entirely within the equatorial plane.
The two axial \(\text{P}-\text{F}\) bond dipoles are positioned directly opposite each other, separated by a \(180^\circ\) angle. Since they have the same magnitude and point in opposite directions, these two opposing forces precisely cancel each other out along the axial direction.
Since the dipoles in both the equatorial plane and along the axial axis cancel, the overall net dipole moment for the entire \(\text{PF}_5\) molecule is zero. This complete cancellation of individual bond polarities is the definitive reason why phosphorus pentafluoride is classified as a nonpolar molecule.