: 200
A molecule's polarity depends on both bond polarity AND geometric symmetry. Dipole moment (mu = q x d) is a vector quantity measured in Debye. Symmetric molecules have zero net dipole despite polar bonds: CO2 (linear), BF3 (trigonal planar), CCl4 (tetrahedral), XeF4 (square planar), SF6 (octahedral). Asymmetric molecules with lone pairs have non-zero dipoles: H2O (1.85 D), NH3 (1.47 D), NF3 (0.24 D), SO2 (1.63 D). The NF3 vs NH3 comparison is a JEE favourite: in NH3, the lone pair dipole and N-H bond dipoles reinforce each other (both point away from H atoms); in NF3, bond dipoles (toward F) oppose the lone pair dipole, nearly cancelling. To determine polarity: draw the structure, identify the geometry, check for symmetry. If all bond dipoles cancel by symmetry and there are no lone pair contributions that break symmetry, the molecule is non-polar. Polarity affects solubility (like dissolves like), boiling point, and intermolecular forces.