VSEPR and Bond Hybridization- AP Chemistry Notes- New Syllabus 2024-2025

Lewis Diagrams and VSEPR Predictions

(i): Molecular Geometry

Molecular geometry is the three-dimensional arrangement of atoms in a molecule. It is determined by the number of regions of electron density (bonding + lone pairs) around the central atom, as predicted by VSEPR theory.

Regions of Electron Density: Each single bond, multiple bond, or lone pair counts as one region.

Basic Electron Domain Geometries:

Molecular Geometry (Adjusted for Lone Pairs):

Key Idea: Molecular geometry results from electron domain geometry modified by lone pairs. Lone pairs push bonding pairs closer together, altering the observed molecular shape.

(ii): Bond Angles

Bond angles are the angles formed between adjacent bonds around a central atom. They are determined by electron domain geometry and modified by lone pairs and multiple bonds due to differences in repulsion strength.

Ideal Bond Angles (no lone pairs):

Lone Pair Effects:

• Lone pairs repel more strongly than bonding pairs.
• This reduces bond angles between bonding atoms compared to ideal values.
• Example: \(\mathrm{NH_3}\) ~107° (vs. 109.5°); \(\mathrm{H_2O}\) ~104.5° (vs. 109.5°).

Multiple Bond Effects:

• Double and triple bonds have greater electron density → repel more strongly than single bonds.
• This can slightly increase angles adjacent to the multiple bond and reduce others.

Key Idea: Bond angles depend on electron domain geometry but are modified by lone pairs and multiple bonds. Lone pairs compress bond angles, while multiple bonds expand angles adjacent to them.

(iii): Relative Bond Energies Based on Bond Order

 Bond energy is the energy required to break a bond in a molecule. It reflects the strength of attraction between bonded atoms. Bond order — the number of shared electron pairs between two atoms — is the primary factor determining bond energy.

Bond Order and Energy:

• Bond order = 1 (single bond) → lowest bond energy.
• Bond order = 2 (double bond) → higher bond energy.
• Bond order = 3 (triple bond) → highest bond energy.

Explanation: Higher bond order means more shared electrons, creating stronger attraction between nuclei and electrons, requiring more energy to break.

Bond Strength Trend: \(\mathrm{single < double < triple}\).

Other Factors: Bond energy also depends on electronegativity difference and atomic size, but bond order is the dominant factor.

Key Idea: Higher bond order increases bond energy because more electrons are shared between atoms, strengthening the attraction and making the bond harder to break.

(iv): Relative Bond Lengths

Bond length is the average distance between the nuclei of two bonded atoms. It depends primarily on bond order (number of electron pairs shared) and the atomic radii of the bonded atoms.

Bond Order and Length:

• Single bond → longest.
• Double bond → shorter.
• Triple bond → shortest.

Reason: Higher bond order increases electron density between nuclei, pulling atoms closer together.

Atomic Radius Effect:

• Bonds between larger atoms are longer because nuclei are farther apart.
• Example: \(\mathrm{C–H}\) bond is shorter than \(\mathrm{C–Cl}\) bond due to smaller H radius.

Bond Length and Strength Relationship:

• Shorter bonds are stronger (higher bond energy).
• Longer bonds are weaker (lower bond energy).

Key Idea: Bond length decreases as bond order increases and increases with larger atomic radii. Shorter bonds are generally stronger than longer bonds.

(v): Presence of a Dipole Moment

A dipole moment occurs when a molecule has an uneven distribution of electron density, resulting in partial positive and negative charges. Dipole moments arise from bond polarity (differences in electronegativity) combined with molecular geometry (vector addition of bond dipoles).

Bond Polarity:

• Bond dipole forms when two atoms have different electronegativities.
• The more electronegative atom gains partial negative charge (\(\delta^-\)), the less electronegative atom gains partial positive charge (\(\delta^+\)).

Molecular Geometry:

• Dipole moments are vectors; molecular geometry determines whether they cancel or add.
• Symmetrical molecules → bond dipoles cancel → nonpolar (e.g., \(\mathrm{CO_2}\)).
• Asymmetrical molecules → bond dipoles reinforce → polar (e.g., \(\mathrm{H_2O}\)).

Dipole Strength:

• Magnitude depends on bond polarity and distance between charges (bond length).
• Measured in Debye units (D).

Key Idea: A molecule has a dipole moment if it contains polar bonds arranged asymmetrically. Bond polarity + geometry together determine molecular polarity.

(vi): Hybridization of Valence Orbitals

Hybridization is the mixing of atomic orbitals (s, p, sometimes d) on the central atom to form new equivalent hybrid orbitals used for bonding. The number of electron domains (regions of electron density) around the central atom determines the type of hybridization and the resulting molecular geometry.

Electron Domains and Hybridization:

Orbital Overlap: Hybrid orbitals form σ bonds with other atoms; unhybridized p orbitals can form π bonds in multiple bonds.

Consistency with VSEPR: Hybridization provides the orbital-level explanation for the bond angles predicted by VSEPR.

Examples:

• \(\mathrm{CH_4}\): sp³ hybridization, tetrahedral.
• \(\mathrm{C_2H_4}\): sp² hybridization (each C), trigonal planar.
• \(\mathrm{C_2H_2}\): sp hybridization (each C), linear.

Key Idea: Hybridization links electron domain geometry (VSEPR) with orbital theory, explaining observed molecular geometries and bond angles at the atomic orbital level.