MCQ Detailed View

The strength of a chemical bond depends on the bond order, orbital overlap, and resonance effects in a molecule. In carbon–carbon bonds, we compare single (C–C), double (C=C), aromatic (delocalized π system), and cyclic bonds.

In CH₃–CH₃ (ethane), the carbon–carbon bond is a pure single bond (sigma bond). A sigma bond is formed by direct head-on overlap of orbitals, which provides maximum strength and stability. The C–C bond in ethane is strong because the overlap between sp³ hybrid orbitals is very effective.

In CH₂=CH₂ (ethene), the carbon–carbon bond is a double bond consisting of one sigma and one pi bond. Although a double bond has a higher bond energy than a single bond, the pi bond is weaker than a sigma bond because of lateral overlap. The overall strength per bond is distributed, and the individual sigma bond strength is not greater than that in ethane.

In benzene, the carbon–carbon bonds are intermediate between a single and double bond due to delocalized π electrons. The bond order in benzene is 1.5, and while delocalization provides stability (resonance energy), the bond strength is not greater than a typical sigma bond.

In cyclohexane, the C–C bonds are single bonds like in ethane. However, due to ring strain and conformational effects, the effective overlap is slightly less than in a free ethane molecule.

Therefore, among the given compounds, the maximum carbon–carbon bond strength is found in CH₃–CH₃ (ethane) because sigma bonds formed by sp³–sp³ overlap provide the strongest C–C interaction in this comparison.