Hybridization Lecture: Part 2

Overview

• Focus on molecules with double and triple bonds.
• Compare with previous discussion on single bonds.

Double Bonds

Example Molecule with Double Bond

• Electronic Geometry: Trigonal planar
• Hybridization: sp²

Carbon Hybridization

• Needs 4 unpaired electrons.
• Hybridization process involves an s orbital and two p orbitals.
• Creates sp² hybrid orbitals and one unhybridized p orbital.

Orbital Overlap

• Blue and red carbon atoms undergo similar hybridization.
• sp² hybrid orbitals overlap to form sigma (σ) bonds.
• Unhybridized p orbitals overlap side-to-side to form a pi (π) bond.
• Sigma Bonds: Stronger, end-to-end overlap
• Pi Bonds: Weaker, side-to-side overlap
• Example: 5 sigma bonds and 1 pi bond in the molecule.

Key Insights

• Double bonds consist of one sigma and one pi bond.
• Bond strength: Double bond not twice the strength of a single bond due to different types of bonds.

Triple Bonds

Example Molecule with Triple Bond

• Electronic Geometry: Linear
• Hybridization: sp

Carbon Hybridization

• Requires 4 unpaired electrons.
• Involves grabbing an empty p orbital for electron pairing.

Orbital Overlap

• Consists of three sigma bonds and two pi bonds.
• Sigma bond: End-to-end overlap of sp orbitals and 1s orbitals of hydrogen.
• Two pi bonds from the side-to-side overlap of remaining p orbitals.

Key Insights

• Triple bond consists of three sigma bonds and two pi bonds.
• Understanding orbital overlaps is crucial for further chemistry studies, particularly in organic chemistry.

Summary

• Sigma Bonds: Found in both single and multiple bonds, result from end-to-end overlap.
• Pi Bonds: Formed by side-to-side overlap, present in double and triple bonds.
• Hybridization affects a molecule's geometry and bonding capabilities.

These notes cover the key concepts regarding the hybridization process, focusing on molecules with double and triple bonds and explaining the differences in bond types and strengths.