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Chapter 5: Molecular Structure and Orbitals

Problems with Lewis Theory

  • Lewis Theory Basics

    • Covalent bonds are pairs of electrons.

  • Major Problems

    • Orbital Role in Covalent Bonding: Doesn't account for orbitals in covalent bonding.

      • Solutions: Hybrid Orbital theory and Molecular Orbital (MO) theory.

    • 3D Shape Prediction: Cannot accurately predict molecule shapes.

      • Solution: Valence Shell Electron Pair Repulsion (VSEPR) theory, which states that bonding electron pairs are maximally spaced due to repulsion.

VSEPR Theory

  • Definition: Valence Shell Electron Pair Repulsion Theory

    • Principle: Arrangement minimizes repulsions among electron domains.

    • Predicts shapes and bond angles in molecules based on electron pair arrangements.

  • Electron Domains: Regions where electron pairs are located.

    • A multiple bond (double or triple) counts as one electron domain.

    • Example: Counting lone pairs as electron domains; four domains include lone pairs.

Electron Domain vs Molecular Geometries

  • Basic Arrangements: 5 common arrangements based on the number of electron groups (2-6).

    • Electron Domain Geometry (EDG): Based on the number of electron domains.

    • Molecular Geometry (MG): Based on the positions of the atoms only.

      • MG differing from EDG due to lone pairs.

Linear Geometry (SN = 2)

  • Shape Characteristics:

    • Even with one bonding and one nonbonding domain, the shape remains linear.

  • Example: BeCl2 is linear.

Trigonal Planar Geometry (SN = 3)

  • Molecular Geometries:

    • Trigonal Planar: All domains bond.

    • Bent: Contains one lone pair.

True Bond Angles: Lone Pairs and Multiple Bonds

  • Nonbonding Pairs: Larger compared to bonding pairs, occupying more space.

  • Effect of Multiple Bonds: Double and triple bonds create more electron density, leading to angle changes due to increased repulsion.

Tetrahedral Geometry (SN = 4)

Electron Groups

Bonding Groups

Lone Pairs

Molecular Geometry

Bond Angles

4

4

0

Tetrahedral

109.5°

4

3

1

Trigonal Pyramidal

<109.5°

4

2

2

Bent

<109.5°

Trigonal Bipyramidal Geometry (SN = 5)

  • Arrangement:

    • Axial and equatorial positions structured such that axial pairs create 90° angles and equatorial pairs create 120° angles.

    • Combination of tetrahedral and trigonal planar configurations.

Octahedral Geometry (SN = 6)

  • Shapes Involving Lone Pairs: Not discussed but pointed towards understanding that they create specific arrangements based on bonding requirements.

VSEPR Summary

  • Predicting Molecular Geometry Steps:

    1. Draw Lewis structure.

    2. Determine the number of electron groups.

    3. Classify as bonding or nonbonding.

    4. Ascertain molecular geometry.

    5. Confirm real bond angles.

Molecular Polarity vs Bond Polarity

  • Polar Bonds: May not result in a polar molecule due to symmetrical geometry canceling out dipoles.

    • Net Dipole Moment: A polar molecule has a nonzero net dipole, typically arising from asymmetrical shapes.

Nonpolar vs Polar Molecules

  • CO2: Two polar bonds create a nonpolar molecule due to linear geometry and equal opposing pulls.

  • H2O: Contains two polar bonds; bent geometry results in a net dipole moment creating a polar molecule.

Bonding: Orbital Overlap

  • Bonds Formed: Achieved through overlapping orbitals, allowing electrons to inhabit a more stable configuration.

Hybridization Basics

  • Definition: Mixing of atomic orbitals creates new hybrid orbitals in chemical bonding.

    • Number of hybrid orbitals equals the number of standard atomic orbitals combined.

  • Hybridization Theories assist in determining molecular shapes and bond characteristics (σ and π bonds).

Hybrid Orbitals (sp, sp2, sp3)

  • sp3 Hybridization: Tetrahedral geometry leading to σ bonds due to head-to-head overlap.

  • sp2 Hybridization: In ethene, gives rise to trigonal planar geometry with one unhybridized p orbital leading to π bonding.

  • Hybrid Orbitals and Structural Arrangements: How hybridization affects molecular characteristics significantly.

Molecular Orbital Theory (MO)

  • Overview: Electrons occupy molecular orbitals similar to atomic ones, where electrons are delocalized and properties like energy levels and sizes are defined.

  • Bonding vs Antibonding: MOs formed from constructive (bonding) vs destructive (antibonding) interactions.

Bond Order

  • Formula: Determines bond strength based on binding vs nonbinding electrons:

    • Bond Order = (Number of bonding e−s - Number of antibonding e−s)/2

    • Relation to bond strength: Higher bond order indicates stronger bonds.

Applications of MO Theory

  • Predictions of Magnetic Properties: Understanding how MO theory predicts properties like magnetism and bond orders in heteronuclear and homonuclear molecules.