Phosphorus tribromide (\(\text{PBr}_3\)) is a common colorless liquid used in organic synthesis, composed of one phosphorus atom bonded to three bromine atoms. Understanding its nature requires examining its internal structure and electron density distribution. The definitive conclusion is that \(\text{PBr}_3\) is a polar molecule, a classification stemming from the differences in atomic electronegativity and the resulting three-dimensional arrangement.
Understanding Molecular Polarity
Molecular polarity describes the overall distribution of electrical charge across a molecule, determined by the presence of polar bonds and the molecule’s shape. This determination relies on electronegativity, which measures an atom’s ability to attract shared electrons in a chemical bond. When atoms with different electronegativity bond, electrons are pulled toward the more electronegative atom, creating a partial negative charge (\(\delta-\)) and a partial positive charge (\(\delta+\)). This unequal sharing forms a polar covalent bond, characterized by a bond dipole moment pointing toward the negative end.
The overall polarity of a molecule depends on the spatial arrangement of these individual bond dipoles, not just the presence of polar bonds. In a symmetrical molecule, the vector sum of all dipoles cancels out, resulting in a net dipole moment of zero and a nonpolar molecule. Conversely, an asymmetrical shape prevents the dipoles from canceling, leaving the molecule with a net, non-zero dipole moment, defining it as polar. Determining polarity requires analyzing both bond polarity and precise three-dimensional geometry.
The Molecular Geometry of \(\text{PBr}_3\)
The structure of \(\text{PBr}_3\) starts with the central phosphorus (P) atom, which has five valence electrons. Three electrons form single covalent bonds with the three surrounding bromine (Br) atoms. This leaves the phosphorus atom with one non-bonding lone pair of electrons. This arrangement results in four distinct regions of electron density around the central atom: three bonding pairs and one lone pair.
To predict the molecular shape, scientists use the Valence Shell Electron Pair Repulsion (VSEPR) theory, which minimizes repulsive forces between electron domains. Based on these four domains, the electron-pair geometry of \(\text{PBr}_3\) is tetrahedral. The molecular geometry, however, differs because the lone pair occupies space but is not an atom. The lone pair exerts increased repulsion on the bonding pairs, forcing the three bromine atoms closer together.
This repulsion compresses the angles between the P-Br bonds to approximately 101 degrees, less than the ideal tetrahedral angle of 109.5 degrees. The resulting shape is called trigonal pyramidal, forming a three-sided pyramid with phosphorus at the apex and the bromine atoms forming the base. The lone pair at the top of the central atom is the source of the molecule’s structural asymmetry.
Why \(\text{PBr}_3\) Is a Polar Molecule
Determining the final polarity of \(\text{PBr}_3\) first requires confirming the polarity of the individual P-Br bonds. Bromine (2.96) is significantly more electronegative than phosphorus (2.19), a difference of 0.77. This confirms that shared electrons are pulled toward the bromine atoms, making each P-Br bond a polar covalent bond.
The second factor is the molecule’s asymmetrical trigonal pyramidal geometry. In this structure, the three bond dipoles point from the less electronegative phosphorus toward the more electronegative bromine atoms, directed downward toward the base of the pyramid. The lone pair on the phosphorus atom also contributes a strong dipole moment directed away from the central atom, compounding the asymmetry.
Since the bond dipoles are oriented in the same general direction and are not opposed by an equal and opposite force, they do not cancel out. This results in a net overall dipole moment for the molecule, experimentally measured at about 0.66 Debye (D). The existence of this measurable net dipole moment definitively classifies \(\text{PBr}_3\) as a polar molecule.