Carbon dioxide (CO_2) has the structure O=C=O in which the central carbon atom forms two double bonds, one with each oxygen atom. Each double bond consists of one sigma (\sigma) bond and one pi (\pi) bond. The sigma bonds arise from the end-to-end (axial) overlap of sp hybrid orbitals of carbon with p orbitals of oxygen, while the pi bonds arise from the sideways (lateral) overlap of unhybridised 2p orbitals. Since CO_2 has two C=O bonds, there are 2 sigma bonds (one per C=O unit) and 2 pi bonds (one per C=O unit). Option '2 sigma, 0 pi' would correspond to a structure with only single bonds (O-C-O), which would not satisfy the octets of oxygen atoms (each O would have only 7 electrons). Option '1 sigma, 2 pi' is incorrect because each double bond always contributes exactly one sigma bond. Option '4 sigma, 0 pi' would imply carbon forms four single bonds, which is not possible with only two oxygen atoms in CO_2 without violating valency rules. Hybridisation: sp carbon uses two sp orbitals for sigma bonds and two unhybridised p orbitals for two pi bonds, giving the linear geometry consistent with the sp hybridisation. This is a direct NCERT hybridisation and bond type question. Plausibility check: CO_2 is a linear molecule and is expected to have two double bonds from a Lewis structure perspective, which exactly gives 2 sigma + 2 pi.