Which of the following is not true about structure of carbon dioxide? (a) In \(\mathrm{CO}_{2}\), carbon is sp - hybridised. (b) C forms two sigma bonds one with each oxygen atom and two \(p \pi-p \pi\) bonds. (c) \(\mathrm{CO}_{2}\) is a linear covalent compound. (d) It is a polar molecule.

Understanding the concept of sp hybridization is crucial when examining the structure of molecules like carbon dioxide (CO_2). Hybridization occurs when atomic orbitals mix to form new, identical hybrid orbitals. In carbon dioxide, the carbon atom has one s orbital and two p orbitals that mix to create two degenerate sp hybrid orbitals. This process involves promoting one of the carbon's 2s electrons to an empty 2p orbital, allowing the formation of two hybridized orbitals.

These new sp orbitals are oriented 180 degrees apart, giving a linear arrangement that is vital for the formation of double bonds with oxygen atoms. Each sp orbital overlaps with an oxygen's p orbital to form a sigma bond, which is the strongest type of covalent bond due to the head-on orbital overlap. The remaining unhybridized p orbitals on carbon are used to form pi bonds with oxygen, which we'll examine in the next section.

Carbon dioxide's double bonds consist of both sigma (σ) and pi (π) bonds. A sigma bond is the first bond formed between two atoms; it is a single bond and allows for free rotation of the bonded atoms around the bond axis. The sigma bonds in CO_2 are formed between the carbon atom's sp hybrid orbitals and the oxygen's p orbitals.

In pi bonds, the overlapping occurs between two parallel unhybridized p orbitals, above and below the plane of the atoms. This second bond is what makes double bonds in molecules like CO_2. Pi bonds are less strong than sigma bonds and restrict the rotation of the bonded atoms, contributing to the molecule's rigidity and shape. One key point to remember is that each double bond is made up of one sigma and one pi bond.

The linear molecular geometry of carbon dioxide is a direct result of the sp hybridization and the sigma and pi bonding we've discussed. When the two sp hybridized orbitals of carbon interact with the p orbitals of the oxygen atoms, they align in a straight line, creating a 180-degree bond angle. This linear arrangement minimizes repulsion between the electron groups around the carbon atom.

In CO_2, this geometry is crucial for understanding various properties of the molecule, including its lack of polarity, as we will explore in the following section. Remember, linear geometry is not unique to CO_2; any molecule with a central atom bonded to two other atoms with no lone pairs will display a similar shape.

Discussions surrounding molecular polarity often reference CO_2 due to its non-polar nature, despite being composed of polar covalent bonds. Polarity is determined by both the electronegativity difference between bonded atoms and molecular geometry. While the bonds between carbon and oxygen are polar because oxygen is more electronegative than carbon, the symmetry of the molecule cancels out any dipole moments.

The linear shape of CO_2 ensures that the pull of the oxygen atoms on the electrons is directly opposed and of equal magnitude, leading to a cancellation of the molecule's net dipole moment. Thus, CO_2 is non-polar, contradicting option (d) in the original exercise. This is an essential concept in chemistry because molecular polarity affects many properties, such as solubility, boiling and melting points, and intermolecular interactions.