Molecular polarity describes the distribution of electric charge across a molecule, influencing its properties and interactions. This charge distribution is quantified by the dipole moment. Whether Phosphorus Trifluoride (\(\text{PF}_3\)) possesses a dipole moment depends entirely on its internal structure and the nature of its chemical bonds.
Understanding Molecular Polarity
A dipole moment measures the separation of positive and negative charges within a molecule. This separation starts at the chemical bond level, known as a bond dipole, which arises from differences in electronegativity. Electronegativity is an atom’s ability to attract shared electrons in a covalent bond.
When two atoms with unequal electronegativity bond, electrons spend more time near the more attractive atom, creating partial negative (\(\delta^-\)) and partial positive (\(\delta^+\)) charges. This unequal sharing establishes a bond dipole, which is a vector quantity with both magnitude and direction.
The presence of polar bonds does not automatically mean the entire molecule is polar. The overall molecular dipole moment is the vector sum of all individual bond dipoles. If the molecular structure is symmetrical, the bond dipole vectors can perfectly cancel, resulting in a net dipole moment of zero. For a molecule to be polar, its structure must be asymmetrical, preventing cancellation and resulting in a measurable net dipole moment.
The Geometry of Phosphorus Trifluoride
Determining the overall polarity of Phosphorus Trifluoride requires establishing its three-dimensional structure. Phosphorus is the central atom, bonded to three Fluorine atoms. The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the spatial arrangement of electron groups around this central atom.
\(\text{PF}_3\) contains 26 valence electrons, resulting in three bonding pairs connecting Phosphorus to the three Fluorine atoms. The remaining two electrons form one non-bonding lone pair, which resides on the central Phosphorus atom. These four electron groups arrange themselves into a tetrahedral electron geometry.
The molecular shape, defined only by the positions of the atoms, is not tetrahedral. The lone pair on the Phosphorus atom pushes the three Fluorine atoms downward, creating a shape known as trigonal pyramidal. This pyramidal arrangement is distinctly asymmetrical, which is the primary factor in determining the molecule’s overall polarity.
The lone pair occupies a greater volume of space than the bonding pairs, distorting the bond angles. The repulsive force of the lone pair compresses the F-P-F bond angles. While the ideal tetrahedral angle is 109.5 degrees, the measured angle in \(\text{PF}_3\) is smaller, approximately 96 to 102 degrees.
Why \(\text{PF}_3\) Possesses a Dipole Moment
The \(\text{P}-\text{F}\) bonds are strongly polar, meeting the first condition for molecular polarity. Fluorine (electronegativity 4.0) is much more electronegative than Phosphorus (electronegativity 2.1). This difference means electron density in each \(\text{P}-\text{F}\) bond is pulled toward the Fluorine atom.
This creates three substantial bond dipoles pointing from Phosphorus toward Fluorine. If \(\text{PF}_3\) were a flat, symmetrical trigonal planar molecule, these equal bond dipoles would cancel each other out, similar to molecules like \(\text{BF}_3\).
The trigonal pyramidal geometry of \(\text{PF}_3\) is inherently asymmetrical. Since the Fluorine atoms are positioned at the base of the pyramid, their individual bond dipole vectors cannot cancel. Instead, the components of the three \(\text{P}-\text{F}\) bond dipoles add together, producing a net dipole moment pointing toward the base of the pyramid.
The non-bonding lone pair of electrons on the Phosphorus atom also contributes to the overall net dipole. This lone pair creates an “orbital dipole” pointing outward from the Phosphorus atom. Both the summed \(\text{P}-\text{F}\) bond dipoles and the lone pair dipole are oriented in the same general direction, reinforcing the charge separation. Because the molecule is asymmetrical and its bond dipoles do not cancel, Phosphorus Trifluoride is a polar molecule and possesses a net dipole moment.