What is the state of hybridization of carbon in CO32- ion?
The Correct Answer and Explanation is:
The carbon in the carbonate ion (CO3^2-) is sp2 hybridized.
Here’s the explanation:
In the CO3^2- ion, the carbon atom is central and is bonded to three oxygen atoms. The overall geometry of the ion is trigonal planar, which suggests that the bond angles around the carbon atom are approximately 120°. This arrangement is typical of sp2 hybridization.
To understand why sp2 hybridization occurs, let’s break down the bonding and electron arrangement:
Bonding and Electron Distribution: The carbonate ion has a resonance structure, meaning the actual bonding is a hybrid of three possible structures where each oxygen atom has a partial negative charge. This suggests that the bonds between carbon and oxygen are not pure single bonds, but a mix of single and partial double bonds (delocalized π bonds). The delocalization of electrons over the three oxygens indicates that the carbon atom’s orbitals overlap in a way that forms three equivalent bonds with oxygen.
Hybridization: For three equivalent bonds, the carbon atom must use three sp2 hybrid orbitals. These orbitals form sigma bonds with the oxygens. The remaining unhybridized p orbital on carbon overlaps with p orbitals on the oxygen atoms to form the delocalized π bonds. This explains the trigonal planar geometry and the bond angles of around 120°, which are characteristic of sp2 hybridization.
Electron Configuration: The carbon atom in CO3^2- has 4 valence electrons. One of these electrons is involved in bonding with an oxygen atom, and the other three are involved in resonance, leading to the delocalization of the electron density across all three C-O bonds.
In conclusion, the carbon atom in the CO3^2- ion undergoes sp2 hybridization, forming a trigonal planar geometry with delocalized π electrons.
