Concept Explanation: Molecular Geometry, Dipole Moment, and Planarity

A molecule has a dipole moment (µ) if there is a separation of charge due to differences in electronegativity and its molecular geometry does not cancel out the individual bond dipoles. A molecule is planar if all its atoms lie in the same plane. For µ = 0, the molecule must be perfectly symmetric so that the vector sum of all its bond dipoles is zero.

Step-by-Step Analysis:

Step 1: Analyze CCl₄ (Carbon Tetrachloride)

Electron Geometry: Tetrahedral (sp³ hybridization).

Molecular Geometry: Tetrahedral.

Symmetry: Highly symmetric. The four C-Cl bond dipoles cancel each other out.

Dipole Moment: µ = 0.

Planarity? No, it is a 3D tetrahedral shape; not all atoms are in one plane.

Conclusion: µ = 0, but not planar.

Step 2: Analyze PCl₅ (Phosphorus Pentachloride)

Electron Geometry: Trigonal Bipyramidal (sp³d hybridization).

Molecular Geometry: Trigonal Bipyramidal.

Symmetry: The axial and equatorial bond dipoles do not cancel perfectly. The two axial P-Cl bonds have a net upward dipole, and the three equatorial bonds have a net dipole in the plane. The vector sum is not zero.

Dipole Moment: µ > 0.

Planarity? No, the axial atoms are out of the plane of the equatorial atoms.

Conclusion: µ ≠ 0, and not planar.

Step 3: Analyze XeF₅^–

Electron Geometry: Octahedral. Xe has 8 valence electrons. In XeF₅^–, it has 5 bonding pairs and 2 lone pairs (one lone pair is on Xe, and the negative charge adds an extra electron, often considered as a lone pair).

Molecular Geometry: Square Pyramidal (if one position is a lone pair).

Symmetry: Not perfectly symmetric. The lone pair creates an asymmetry, and the bond dipoles do not cancel.

Dipole Moment: µ > 0.

Planarity? The five F atoms are not in one plane; one is axial in the pyramid.

Conclusion: µ ≠ 0, and not planar.

Step 4: Analyze PCl₄⁺

Electron Geometry: Tetrahedral. P has 5 valence electrons, but in the cation PCl₄⁺, it loses one electron, leaving 4 valence electrons for bonding (no lone pairs).

Molecular Geometry: Tetrahedral.

Symmetry: Perfectly symmetric, like CCl₄. The four equivalent P-Cl bond dipoles cancel each other out.

Dipole Moment: µ = 0.

Planarity? No, it is tetrahedral and not planar.

Conclusion: µ = 0, but not planar.

Final Answer: None of the given options are both planar and have µ = 0. CCl₄ and PCl₄⁺ have µ = 0 but are not planar. PCl₅ and XeF₅^– are neither planar nor have µ = 0. The question might contain a trick, or there could be a misinterpretation. Based on standard knowledge, no option fully satisfies both conditions.

Related Topics:

• VSEPR Theory: Predicts molecular geometry based on electron pairs around the central atom.
• Hybridization: Explains the mixing of atomic orbitals to form new hybrid orbitals for bonding.
• Symmetry and Point Groups: Used to determine molecular properties like dipole moment.
• Bond Polarity and Dipole Moments: Arise from electronegativity differences between atoms.

Key Formulae and Theory:

Dipole Moment (µ): $\mu =Q\times r$, where Q is the magnitude of charge separation and r is the distance between charges.

Vector Sum: The net dipole moment is the vector sum of all individual bond moments. For µ to be zero, the molecule must have a symmetric structure where these vectors cancel out.

Common Geometries with µ = 0: Linear (e.g., CO₂), Tetrahedral (e.g., CH₄, CCl₄), Octahedral (e.g., SF₆), Square Planar (e.g., XeF₄).

Planar Molecules: Examples include BF₃ (trigonal planar), C₂H₄ (trigonal planar around C atoms), XeF₄ (square planar).