Q.Which of the following is true for CO2?
[Table — option / nature of the C=O bond / nature of the CO2 molecule]
(A) polar / non-polar
(B) non-polar / polar
(C) polar / polar
(D) non-polar / non-polar

A covalent bond's shared electron pair is shared perfectly equally only when the two bonded atoms have the same or very similar electronegativity (e.g. H–H, F–F) — a nonpolar covalent bond, with the shared pair sitting exactly midway between the nuclei. Where the two atoms differ noticeably in electronegativity (e.g. H–F), the shared pair is pulled closer to the more electronegative atom, creating a dipole: a partial negative charge (δ−) on that atom and a partial positive charge (δ+) on the other — a polar covalent bond, whose polarity increases as the electronegativity difference increases (about 1.7 difference ≈ 50% ionic character). Dipole moment ($\mu$) is formally defined as the product of the separated charge and its separation distance, $\mu = Q\times r$, a VECTOR quantity conventionally drawn as a small arrow from the positive centre to the negative centre, measured in Debye (D), where $1\ \text{D} = 3.33564\times10^{-30}\ \text{C m}$. For a POLYATOMIC molecule, the overall (net) dipole moment is the VECTOR SUM of every individual bond dipole plus any lone-pair contribution, so it depends on both the molecule's shape and the size/direction of each contributing dipole — which is why some molecules built entirely from polar bonds still have ZERO net dipole moment when their shape is symmetric enough for the individual bond dipoles to cancel exactly (linear BeF2; trigon …