Molecular polarity describes the fundamental way electrical charge is distributed across a chemical compound. If the charge is distributed unevenly, the molecule possesses a positive and a negative end, resulting in polarity. Phosphorus Oxychloride, \(\text{POCl}_3\), is definitively a polar molecule. This polarity results directly from the chemical bonds within the structure and the three-dimensional arrangement of its atoms.
The Building Blocks: Understanding Chemical Bonds
Determining overall molecular polarity begins by analyzing the polarity of the individual bonds. Bond polarity is governed by electronegativity, an atom’s ability to attract shared electrons. When atoms with different electronegativity values bond, the shared electrons spend more time near the more attractive atom, creating partial positive and negative charges.
In \(\text{POCl}_3\), the central phosphorus (P) atom bonds to one oxygen (O) atom and three chlorine (Cl) atoms. Oxygen (3.44) and chlorine (3.16) are significantly more electronegative than phosphorus (2.19). This difference ensures that both the phosphorus-oxygen (P-O) and phosphorus-chlorine (P-Cl) bonds are polar covalent bonds.
The P-O bond has a larger electronegativity difference (1.25) compared to the P-Cl bonds (0.97), making the P-O bond the most polar of the four. Electron density is pulled strongly toward the oxygen atom in the P-O bond and toward the chlorine atoms in the P-Cl bonds.
Determining Molecular Shape
Overall polarity depends not just on polar bonds, but also on their spatial arrangement. The arrangement of atoms in \(\text{POCl}_3\) is predicted using the Valence Shell Electron Pair Repulsion (VSEPR) theory. VSEPR theory states that electron domains arrange themselves around the central atom to be as far apart as possible to minimize repulsion.
The central phosphorus atom forms four bonding domains: a double bond to oxygen and three single bonds to chlorine atoms. Since there are no lone pairs on the phosphorus atom, the four domains adopt a tetrahedral electron geometry. This arrangement results in a molecular shape described as a distorted tetrahedron.
The molecular shape is inherently asymmetrical because the four surrounding atoms are not identical. The oxygen atom is chemically distinct from the three chlorine atoms, preventing the molecule from achieving perfect symmetry. This lack of symmetry is necessary for polarity, as it ensures the electrical forces generated by the individual bonds will not be balanced.
Vector Summation and Net Dipole Moment
The final determination of molecular polarity requires combining the information about bond polarity and molecular shape through vector summation. Each polar bond acts like a vector, pointing from the less electronegative atom toward the more electronegative atom. The overall molecular polarity is the net result of adding these four bond vectors together.
In a perfectly symmetrical molecule, such as carbon tetrachloride (\(\text{CCl}_4\)), the bond vectors are equal in magnitude and point in opposite directions, causing them to cancel each other out and result in a nonpolar molecule. However, for \(\text{POCl}_3\), the P-O bond vector is significantly different from the three P-Cl bond vectors, both in magnitude and direction relative to the others.
The strong pull of the highly electronegative oxygen atom creates a large dipole moment. Because the molecular shape is asymmetrical, the four individual bond vectors cannot cancel out. This imbalance leads to a net dipole moment, confirming \(\text{POCl}_3\) as a polar molecule with a permanent separation of charge.