Chemical Bonding II: Advanced Concepts

Chemical Bonding II: Advanced Concepts

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 Valence Bond (VB) theory.

  4. Predict the hybridization of an atom.

  5. Distinguish between sigma (σ) and pi (π) bonds.

  6. Identify the atomic and/or hybrid orbitals used to make a particular bond.

  7. Distinguish between bonding and anti-bonding molecular orbitals (MOs).

  8. Draw and interpret MO diagrams for simple diatomic molecules.

  9. Use MO theory to predict properties of diatomic molecules such as bond order, bond strength, and magnetism.

Lewis Structure of Water

  • The Lewis theory, developed in 1916, predicts that there are regions of electrons in an atom.

  • Regions can result from placing shared pairs of valence electrons between bonding nuclei or from placing unshared valence electrons on a single nucleus.

  • Lewis structures predict:

    • a linear (180°) shape for diatomic molecules.

    • a right angle (90°) shape for some polyatomic molecules.

  • The bond angle in water (H₂O) is between these extremes due to electron group repulsion aiming to minimize interactions.

Predicting Molecular Geometry

  • Molecular geometry and the number of valence electron pairs should be correlated.

  • Introduced in 1939, Valence-Shell Electron-Pair Repulsion (VSEPR) Theory further explains this correlation.

    • Proposed independently by Ryutaro Tsuchida in 1939, and later developed by Nevil Sidgwick and Herbert Powell in 1940, and further refined by Ronald Gillespie and Ronald Sydney Nyholm in 1957.

VSEPR Theory

  • VSEPR Theory extends Lewis theory.

  • It predicts molecular geometry by arranging electron pairs (groups) to minimize electrostatic repulsions between them.

  • There are five basic molecular shapes derived from VSEPR theory:

    1. Linear

    2. Trigonal planar

    3. Tetrahedral

    4. Trigonal bipyramidal

    5. Octahedral

Electron Group vs. Molecular Geometry

  • Electron Group Geometry: refers to the spatial arrangement of all electron groups (bonding and lone pairs) around a central atom, which follows the five basic shapes.

  • Molecular Geometry: refers to the actual arrangement of atoms in space, defined by the connections between the atomic nuclei.

    • Example: Electron group geometry may be tetrahedral while the molecular geometry of water is bent, caused by lone pair repulsion.

VSEPR Designation

  • Molecules or polyatomic ions are assigned an AXmEn designation:

    • A: Central atom

    • X: Bonded atoms

    • E: Nonbonding valence electron groups (usually lone pairs)

    • m & n: Integers representing the number of X and E groups, respectively.

VSEPR Geometries

Table 10.1: Electron and Molecular Geometries

    • Electron Groups

    Bonding Groups

    Lone Pairs

    Electronic Geometry

    Molecular Geometry

    Approximate Bond Angles

    3

    3

    0

    Trigonal planar

    Trigonal planar

    120°

    3

    2

    1

    Trigonal planar

    Bent

    <120°

    4

    4

    0

    Tetrahedral

    Tetrahedral

    109.5°

    4

    3

    1

    Tetrahedral

    Trigonal pyramidal

    <109.5°

    4

    2

    2

    Tetrahedral

    Bent

    <109.5°

    Drawing 3D Molecules on 2D Paper

    • To represent 3D shapes on 2D paper:

      • The central atom is placed in the plane of the paper.

      • Atoms in the same plane are indicated with a straight line.

      • For atoms in front of the plane, use a solid wedge; for those behind the plane, use a hashed wedge.

    Examples and Exercises

    • Example Exercise: Draw the structure of CO₃²⁻ using wedges and dashes.

    • Your Turn: For CO₃²⁻, identify the molecular shape:

      • a) Trigonal planar

      • b) Trigonal pyramidal

      • c) Bent

      • d) Tetrahedral

      • e) Trigonal bipyramidal

    Electric Geometry and Molecular Geometry of Water

    • Discusses how lone-pair versus bonding-pair repulsions influence overall molecular geometry predictions.

      • Lone pairs exert greater repulsive forces compared to bonded pairs, affecting the angles between bonded atoms.

    Identifying Molecular Geometry

    • Lewis Structure should be drawn first.

      • Verify the electron groups around the central atom.

      • Count the lone pairs and determine molecular geometry.

    Intermolecular Forces (IMFs)

    • IMFs refer to the forces holding molecules together in a liquid or solid.

    • IMFs are significantly weaker than covalent bonds and play a crucial role in determining bulk properties like melting and boiling points.

    Types of Intermolecular Forces

    1. London Dispersion Forces (van der Waals forces):

      • Occur in nonpolar molecules due to temporary fluctuations in electron distributions, leading to induced dipole moments.

    2. Dipole-Dipole Interactions:

      • Occur between polar molecules with permanent dipole moments, resulting in higher boiling and melting points relative to nonpolar molecules.

    3. Hydrogen Bonding:

      • Strong attractions occur between a hydrogen atom covalently bonded to highly electronegative atoms (O, N, F), resulting in notably high boiling points.

    4. Ion-Dipole Interactions:

      • Occur in solutions where ions dissociate and interact with polar molecules, crucial for solubility of ionic compounds in polar solvents.

    Molecular Polarity

    • For a molecule to be polar:

      1. It must have polar bonds with an electronegativity difference between the atoms.

      2. The shape must be unsymmetrical to yield a net dipole.

      3. Nonbonding pairs can influence molecular polarity, producing a strong effect on the overall dipole.

    Homework Assignments

    • Determine molecular geometries and sketch for various molecules including SF₄, ClF₃, IF₂⁻, and IBr₄⁻.

    • Write hybridization and bonding schemes for compounds like COCl₂ and NH₂⁻, including identifying sigma and pi bonds.

    Comparison of Bonding Theories

    • Lewis Bonding Theory:

      • Easy and quick for predicting covalent and ionic bonds but does not account for molecular shape.

    • VSEPR Theory:

      • Explains shapes based on electrostatic interactions.

    • Valence Bond Theory:

      • Integrates Lewis & VSEPR to explain bond formation as the overlap of atomic orbitals.

    • Molecular Orbital Theory (MOT):

      • Describes electrons in a molecule being delocalized across the entire molecule, providing a comprehensive view of bonding.

    Conclusion

    • Molecular geometry, hybridization, and investigation of intermolecular forces provide a comprehensive framework for understanding chemical bonding and molecular interactions, essential for predicting molecular behavior in practical applications.