Question

C-Cl bond in methyl chloride compared to C-Cl bond in chlorobenzene is

• A. Longer and stronger
• B. Shorter and stronger
• C. Shorter and weaker
• D. Longer and weaker

Hint:

Remember these two key factors when comparing bonds to a benzene ring with bonds to a simple alkyl group: 1. \textbf{Hybridization:} A bond to an sp\(^2\) carbon is shorter and stronger than a bond to an sp\(^3\) carbon. 2. \textbf{Resonance:} Resonance can introduce partial double bond character, which shortens and strengthens the bond. Both factors work in the same direction for chlorobenzene, making its C-Cl bond shorter and stronger.

Answer 1

The Correct Option is D

Step 1: Understanding the Concept:
Bond length and bond strength are primarily dictated by the state of hybridization of the bonded carbon atom and the presence or absence of electron delocalization (resonance).
Step 2: Key Formula or Approach:
Compare the hybridization ($sp^3$ vs $sp^2$) of the carbon atoms attached to chlorine in both molecules and analyze the partial double bond character arising from resonance in aromatic systems.
Step 3: Detailed Explanation:
1. In Methyl Chloride ($CH_{3}Cl$): The carbon atom is $sp^{3}$ hybridized. The $C-Cl$ bond is a pure, localized single sigma bond.

2. In Chlorobenzene ($C_{6}H_{5}Cl$):
- The carbon atom attached to the chlorine is $sp^{2}$ hybridized. $sp^{2}$ hybridized carbons have more s-character, making them smaller and more electronegative, which inherently pulls the chlorine atom closer.
- More importantly, Resonance: The lone pairs of electrons on the Chlorine atom are delocalized into the $\pi$ system of the benzene ring. This imparts a partial double bond character to the $C-Cl$ bond.
Because double bonds are shorter and stronger than single bonds, the $C-Cl$ bond in chlorobenzene is shorter and stronger.
Therefore, by relative comparison, the pure single bond in methyl chloride is longer and weaker.
Step 4: Final Answer:
The C-Cl bond in methyl chloride is longer and weaker.