A molecule's polarity is determined by the vector sum of its individual bond dipoles .

1. $SbCl_5$: The central antimony atom ($Sb$) undergoes $sp^3d$ hybridisation, resulting in a trigonal bipyramidal geometry . Because all five terminal atoms are identical (chlorine), the three equatorial bond dipoles cancel each other at $120^\circ$, and the two axial bond dipoles cancel each other at $180^\circ$ . Consequently, the net dipole moment ($\mu$) is zero, making it non-polar.
2. $NO_2$: This is an odd-electron molecule with a bent geometry . The asymmetrical arrangement prevents bond dipoles from cancelling, resulting in a polar molecule .
3. $POCl_3$: The central phosphorus atom is $sp^3$ hybridised (tetrahedral-like). Since the $P=O$ bond dipole differs in magnitude and direction from the three $P-Cl$ bond dipoles, they do not cancel out, making it polar.
4. $CH_2O$ (Formaldehyde): The carbonyl carbon is $sp^2$ hybridised . The $C=O$ double bond is highly polar due to the higher electronegativity of oxygen relative to carbon . The geometry is trigonal planar, but the different dipoles of $C-H$ and $C=O$ bonds do not cancel, resulting in a significant net dipole moment.