Phosphorus pentachloride (\(\text{PCl}_5\)) is an inorganic compound utilized in chemical synthesis as a powerful chlorinating agent. This compound is formed when one central phosphorus atom bonds with five chlorine atoms. Despite having bonds with unequal electron sharing, the entire \(\text{PCl}_5\) molecule is nonpolar. Understanding this lack of overall polarity requires examining how its internal structure influences charge distribution.
Defining Molecular Polarity
Molecular polarity describes the uneven sharing of electrons between atoms, which creates a separation of charge. This concept is based on electronegativity, which is an atom’s ability to attract electrons in a chemical bond. When atoms of differing electronegativity bond, electrons spend more time near the more attractive atom, resulting in a polar covalent bond. This unequal sharing generates a bond dipole.
The atom with greater electron density develops a partial negative charge, and the other atom develops a partial positive charge. For a molecule to be polar, it must possess a net dipole moment, meaning the vector sum of all individual bond dipoles is greater than zero. A molecule is nonpolar if its net dipole moment is zero, even if its bonds are polar. This cancellation happens when bond dipoles are arranged symmetrically, pulling in opposite and equal directions. Therefore, polarity depends on both bond nature and three-dimensional shape.
The Three-Dimensional Structure of Phosphorus Pentachloride
To determine the three-dimensional arrangement of atoms in \(\text{PCl}_5\), scientists apply the Valence Shell Electron Pair Repulsion (VSEPR) theory. This theory predicts molecular geometry based on the principle that electron domains around a central atom arrange themselves to minimize repulsion. In \(\text{PCl}_5\), the central phosphorus (P) atom is bonded to five chlorine (Cl) atoms and has no lone pairs.
The presence of five bonding domains results in the geometric arrangement known as trigonal bipyramidal. This shape is characterized by two distinct sets of positions for the chlorine atoms. Three chlorine atoms occupy equatorial positions, lying in a triangular plane around the phosphorus center with bond angles of \(120^\circ\). The remaining two chlorine atoms occupy axial positions, located directly above and below the equatorial plane. These two axial bonds are oriented \(90^\circ\) to the equatorial plane.
Symmetry and Net Dipole Moment
The nonpolar nature of \(\text{PCl}_5\) is a direct consequence of its trigonal bipyramidal geometry and symmetry. Although the phosphorus-chlorine (\(\text{P-Cl}\)) bonds are individually polar because chlorine is more electronegative than phosphorus, the symmetrical arrangement ensures the bond dipoles cancel each other out.
The three \(\text{P-Cl}\) bonds in the equatorial plane are separated by \(120^\circ\). The dipole moment created by the electron pull toward one equatorial chlorine atom is exactly balanced by the combined pull of the other two. These three vectors sum to zero, meaning the equatorial plane contributes no net dipole moment. Simultaneously, the two \(\text{P-Cl}\) bonds in the axial positions are oriented \(180^\circ\) from one another, pulling electrons in exactly opposite directions. The dipole from the upper axial chlorine is cancelled out by the dipole from the lower axial chlorine. Since the net dipole moment of both the equatorial plane and the axial axis is zero, the overall molecule has a net dipole moment of zero.