The molecule PCl5 is observed not to have a dipole moment. This is because:
A. There are no lone pairs of electrons on the central atom
B. As a gas and liquid PCl5 is not a ionic, but rather the bonds are covalent
C. The polarity of the P-Cl bonds cancel out due to the geometry of the molecule.
D. There are two lone pairs of electrons on the central atom, but due to the repulsion, they are on opposite sides of the central atom and cancel out.
E. P and Cl are ciose in the periodic table, so they have very similar electronegatives, and as such, the P-Cl bonds are not polar.
The correct answer and explanation is :
The correct answer is:
C. The polarity of the P-Cl bonds cancel out due to the geometry of the molecule.
Explanation:
Phosphorus pentachloride (PCl₅) is a molecule where phosphorus (P) is the central atom bonded to five chlorine (Cl) atoms. To understand why PCl₅ has no dipole moment, we need to consider both the bond polarity and the molecular geometry.
- Bond Polarity:
Each P–Cl bond is polar because phosphorus and chlorine differ in electronegativity (chlorine is more electronegative than phosphorus). This difference causes a partial negative charge (δ⁻) on chlorine and a partial positive charge (δ⁺) on phosphorus in each bond, creating a dipole moment along each P–Cl bond.
- Molecular Geometry:
The shape of PCl₅ is trigonal bipyramidal, based on the Valence Shell Electron Pair Repulsion (VSEPR) theory. This geometry consists of:
- Three chlorine atoms arranged in an equatorial plane at 120° angles.
- Two chlorine atoms positioned axially, 180° apart and perpendicular to the equatorial plane.
- Dipole Moment Cancellation:
Despite the polarity of individual bonds, the overall dipole moment depends on how these bond dipoles add up vectorially.
- The three equatorial P–Cl bond dipoles are arranged symmetrically and cancel each other out because their vectors sum to zero.
- The two axial P–Cl bond dipoles are equal in magnitude but opposite in direction (180° apart), so they cancel each other out as well.
When all bond dipoles are combined, the symmetrical trigonal bipyramidal geometry causes the dipole vectors to cancel, leading to an overall net dipole moment of zero.
Why the other options are incorrect:
- A: The absence of lone pairs alone does not guarantee no dipole moment. For example, molecules like NH₃ have no lone pairs on the central atom but are polar due to shape.
- B: Whether PCl₅ is ionic or covalent is irrelevant to the net dipole moment; the dipole depends on molecular shape and bond polarity.
- D: PCl₅ does not have lone pairs on phosphorus in its common form. It uses all five valence electrons to bond with chlorine atoms.
- E: Although P and Cl are neighbors in the periodic table, they still have a sufficient electronegativity difference to make the bonds polar; the lack of dipole moment is due to geometry, not lack of bond polarity.
In summary, PCl₅ has polar bonds but a symmetrical trigonal bipyramidal geometry, which causes the individual bond dipoles to cancel out, resulting in a molecule with no overall dipole moment.