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Chemical Bonding II: Advanced Concepts

Chemical Bonding II: Advanced Concepts Notes

Module Learning Goals

By the end of this module, students will be able to:

  1. Predict the electron group geometry and/or molecular geometry of a molecule using VSEPR theory.
  2. Predict whether a molecule will have a net dipole moment.
  3. Describe the bonding in small molecules using VB theory.
  4. Predict the hybridization of an atom.
  5. Distinguish between sigma and pi bonds.
  6. Identify the atomic and/or hybrid orbitals used to make a particular bond.
  7. Distinguish between bonding and anti-bonding MOs.
  8. Draw and interpret MO diagrams for simple diatomics.
  9. Use MO theory to predict properties of diatomic molecules (bond order, bond strength, magnetism).

Lewis Structure of Water

  • Lewis theory developed in 1916 predicts that there are regions of electrons in an atom.
  • Some of these regions arise from placing shared pairs of valence electrons between bonding nuclei.
  • Others arise from placing unshared valence electrons on a single nucleus.
  • Lewis structures predict:
    • A linear (180°) or right angle (90°) shape for water.
    • The actual H-O-H bond angle is between these, as electron groups strive to minimize their interaction.

Predicting Molecular Geometry

  • There is a correlation between molecular geometry and the number of valence electron pairs, defined by Valence-Shell Electron-Pair Repulsion (VSEPR) Theory.
  • VSEPR Theory:
    • Contribution Timeline:
    • Original proposal in 1939 by Ryutaro Tsuchida.
    • Similar ideas presented by Nevil Sidgwick and Herbert Powell in 1940.
    • Further-developed by Ronald Gillespie and Ronald Sydney Nyholm in 1957.
    • Predicts molecular geometry by arranging electron pairs to minimize electrostatic repulsions.

Electron Group vs. Molecular Geometry

  • Electron Group Geometry:
    • The geometrical distribution of electron groups (e.g., bond and lone pairs).
    • Follows five basic shapes.
  • Molecular Geometry:
    • The geometrical arrangements of atomic nuclei.
    • For example:
    • Electron group geometry: tetrahedral
    • Molecular geometry: bent

Double and Triple Bonds

  • VSEPR theory treats all bonds equally (single, double, etc.) and views a bond as an electron group.
  • Examples:
    • 2 electron groups yield linear geometry.
    • 3 electron groups yield 120° trigonal planar geometry.

AXmEn Designation with VSEPR

  • A molecule or polyatomic ion is assigned an AXmEn designation, where:
    • A = central atom
    • X = bonded atom
    • E = nonbonding valence electron group (lone pair)
    • m and n are integers representing the number of bonded atoms and lone pairs, respectively.
  • Each group around the central atom is designated either a bonding group (BG) or a lone pair (LP).

VSEPR Geometries

  • Table 10.1 summarizes the VSEPR geometries:
    • Electron Geometry:
    • Trigonal planar: 3 bonding groups, 0 lone pairs, 120° bond angles.
    • Tetrahedral: 4 bonding groups, 0 lone pairs, 109.5° bond angles.
    • Trigonal pyramidal: 4 groups, 3 bonding, 1 lone pair, <109.5° bond angles.
    • Bent: 4 groups, 2 bonding, 2 lone pairs, <109.5° bond angles.

Drawing Three-Dimensional Molecules

  • Molecules can be represented in three dimensions on a two-dimensional plane using wedges and dashes:
    • Solid Wedge: Atom is in front of the plane.
    • Dashed Wedge: Atom is behind the plane.

Lone Pairs and Molecular Geometry

  • Lone pairs exert greater repulsion than bonding pairs, leading to adjustments in bond angles.
  • Observations:
    • CH4 (4 bonding pairs): tetrahedral geometry
    • NH3 (3 bonding pairs, 1 lone pair): trigonal pyramidal geometry <109.5°
    • H2O (2 bonding pairs, 2 lone pairs): bent geometry <109.5°

Applying VSEPR Theory

  1. Draw a Lewis Structure.
  2. Count the number of electron groups around the central atom.
  3. Determine the electron group geometry.
  4. Count the number of lone pairs on the central atom.
  5. Determine the molecular geometry.

Molecular Shape and Polarity

  • For a molecule to be polar:
    • Must have polar bonds (electronegativity difference).
    • Must exhibit an unsymmetrical shape.
  • The net dipole moment is calculated using vector addition of individual dipole moments within the molecule.

Dipole Moments and Molecular Geometry

  • Polar bonds create dipole moments, which are vector quantities with magnitude and direction.
  • Molecular dipole moments depend on individual bond dipole moments and overall molecular geometry.

Intermolecular Forces

  • Intermolecular forces are attractive forces holding molecules together in liquids/solids and are weaker than covalent bonds.
  • Types of Intermolecular Forces:
    1. London Dispersion Forces (Van der Waals forces): Present in all molecules, temporary fluctuations in electron distribution.
    2. Dipole-Dipole Interactions: Permanent dipoles in polar molecules.
    3. Hydrogen Bonding: Strong interactions between H and electronegative atoms (O, N, F).
    4. Ion-Dipole Attraction: Occurs in mixtures of ionic and polar compounds.

Hydrogen Bonding

  • Strong interactions resulting from large electronegativity differences in H-O, H-N, H-F bonds.
  • Leads to unique properties of substances like water (high boiling point, solvent properties).

Solubility and Attractive Forces

  • Solubility depends on attractive forces between solute and solvent molecules (like dissolves like).
  • Polar substances dissolve in polar solvents, ionic compounds are soluble in water due to ion-dipole interactions, while nonpolar substances dissolve in nonpolar solvents.

Problems with VSEPR and Lewis Theory

  • Lewis structures assume all bonds are equivalent and cannot accurately predict molecular properties (e.g., magnetism of O2).
  • Valence Bond Theory (VBT): Bonds are formed through overlapping atomic orbitals; predicts how sigma and pi bonds are formed.

Molecular Orbital Theory (MOT)

  • Provides a more thorough description of bonding by accounting for geometry, magnetic properties, and violations of the octet rule.
  • Involves the linear combination of atomic orbitals (LCAO) to describe molecular orbitals.

Key Concepts in Molecular Orbital Theory

  • Electrons fill molecular orbitals from lowest to highest energy.
  • Bonding molecular orbitals are lower in energy while antibonding orbitals are higher in energy.
  • A stable molecule has more electrons in bonding MOs than in antibonding MOs (bond order >0).

Bond Order in MO Theory

  • Bond Order = (Number of electrons in bonding MOs - Number of electrons in antibonding MOs) / 2.

Molecular Orbital Diagrams

  • Used to represent molecular orbitals, showing the relative energy levels of bonding and antibonding orbitals.

Summary of Theories

  • Lewis Bonding Theory: Simple and quick, does not predict shapes.
  • VSEPR Theory: Helps predict molecular shapes based on electrostatic repulsion.
  • Valence Bond Theory: Incorporates quantum mechanics for better predictions based on orbital overlaps.
  • Molecular Orbital Theory: Offers a complete description of bonding and molecular properties, resolving issues found in earlier theories.

Laboratory Applications

  • Students apply these theories with practical examples and molecular modeling exercises to reinforce learning outcomes.