lecture 14: Hybridization, Resonance, and Peptide Bonds

Lecture 14 - 09/29/2025

Hybridization

• Methanol (CH₃OH)

  • Structure:

  • H–C–O–H

  • Angle 1: (C–H–C) = tetrahedral = $109.5^ ext{o}$

  • Angle 2: (C–O–H) = bent = _

  • Hybridization of C: sp³ single bonds

• Electron Density:

  • Regions of Electron Density: Only 3 regions, including double bond.

  • Electron Geometry: Repulsion places them at the corners of a trigonal planar shape, leading to 120° bond angles.

• Hybridization Types:

  • sp² Hybridization: C with one double bond (e.g., formaldehyde, CH₂O)

  • Structure:

    • H
      • H–C–O
      •

  • Hybridization of C: sp² with 1 double bond

Degrees of Hybridization

• 2 Electron Domains:

  • Electron Geometry = linear

  • Hybridization of C = sp with 2 double bonds

Resonance and Hybridization

• Glycine (Amino Acid):

  • Resonance affects hybridization, indicating atoms adjacent to double bonds are often sp² hybridized.

• Carboxylic Acids and Esters:

  • Carbon and both oxygens are sp² hybridized due to resonance.

Peptide Bonds

• Peptide Bond Specifications:

  • R' and R groups attached to nitrogen (N) and carbon (C).

  • Partial Double Bond: The nature of bonding leads to trigonal planar geometries.

  • Hybridization:

  • N and C are each sp² hybridized leading to planar geometry.

Hybridization Properties of Carbon

1. Octet Completion: Carbon completes its octet by sharing electrons.

2. Electron Configuration: [He]2s²2p², which makes ion formation difficult (would have to gain or lose 4 electrons).

3. Electronegativity: Intermediate electronegativity of 2.5.

4. Covalent Bonds: Carbon typically forms four covalent bonds.

5. Types of Hybridization: Carbon can be hybridized as sp³, sp², and sp.

6. Bond Formation: Carbon can form single, double, and triple bonds.

Orbital Hybridization

Bonds Explanation

• Sigma (σ) Bonds:

  • Defined as single bonds where electrons are shared between atoms with an overlap region lying directly between the two nuclei and concentrated along the internuclear axis.

  • Each carbon typically has four electron domains (e.g., in ethane (C₂H₆)).

  • Therefore, expect sp³ hybridization with no remaining p orbitals for sigma bonds.

Hybridization in Molecules with Double Bonds

• Ethylene (C₂H₄): Contains five sigma bonds involving carbon-carbon and carbon-hydrogen bonds.

• Unhybridized 2p orbitals are involved in pi (π) bonding, which is critical for double bonds.

• Ethylene structure: H–C=C–H

• The unhybridized 2p orbitals allow for sideways overlaps to form π bonds.

• Sigma and Pi Bonds in Ethylene:

• A sigma bond is formed by the overlap of sp² hybrid orbitals, while the π bond is formed by the overlap of p orbitals.

Acetylene (C₂H₂) Hybridization Characteristics

• Bonding:

  • 3 sigma bonds (1 C-C and 2 C-H)

  • 2 pi bonds formed between C-C from unhybridized p orbitals.

Bonding in CH₂O and CH₃COOH (Acetic Acid)

• CH₂O

  • Bonding Characteristics:

  • Carbon: 3 bonded atoms, no lone pairs, hybridization sp².

  • Oxygen: sp² hybridization leading to 3 sigma and 1 pi bonds, where the C=O double bond consists of 1 σ bond and 1 π bond.

• Acetic Acid (CH₃COOH):

  • Total Bonds:

  • Sigma bonds = 6 + 1 = 7

  • Pi bonds = 1

Consequences of Multiple Bonding

• Rotation Restrictions:

  • Rotation around double & triple bonds is severely restricted.

  • Sigma bonds (single bonds) allow for free rotation.

Molecular Geometry and Polarity

• Isomer Distinction: Dipole moments distinguish structural isomers.

• Cis-Trans Isomerism: Double bonds prevent rotation, leading to distinct geometric configurations.

  • The cis isomer is polar while the trans isomer is nonpolar.

  • Boiling point of the cis isomer is 13°C higher than that of the trans isomer due to polarity differences.