Final answer:
In CO, the carbon atom is sp hybridized; in CO2, the carbon atom is also sp hybridized; and in CN-, the carbon atom is sp hybridized. These hybridizations correspond to the number and types of bonds each carbon forms and their molecular geometries, which are linear for all these species.
Step-by-step explanation:
The hybrid orbitals used by carbon atoms in the species CO, CO₂, and CN⁻ can be identified based on their molecular structures and the bond types they contain. For each species, the following hybridizations are present:
- CO (carbon monoxide) has a carbon atom that is sp hybridized because it forms two sigma bonds and has two pi bonds, which corresponds to a linear molecular geometry.
- CO₂ (carbon dioxide) carbon atom is also sp hybridized, as it forms two double bonds (each consisting of one sigma and one pi bond) and the molecule has a linear shape.
- CN⁻ (cyanide ion) has a carbon atom that is sp hybridized, indicated by the triple bond (consisting of one sigma and two pi bonds) to nitrogen and one single bond, resulting in a linear geometry.
It is important to note that sigma bonds are typically formed by the end-to-end overlap of hybrid orbitals, whereas pi bonds result from the side-by-side overlap of unhybridized p orbitals. Therefore, the hybrid orbitals are used to form sigma bonds, and the remaining unhybridized p orbitals form pi bonds in molecules with multiple bonds.