Chemical Bonding II: Molecular Shapes

Chemical Bonding II: Molecular Shapes

1. VSEPR Theory: The 5 Basic Shapes & the Effect of Lone Pairs

• Bond Length:
  • Refers to the distance between nuclei (denoted as d1, d2).
• Bond Angle:
  • Defined as the angle between adjacent bonds (denoted as α).

2. VSEPR Model (Valence Shell Electron Pair Repulsion)

• Definition: A molecular shape prediction model.
• Main Postulate:
  • The structure surrounding a given atom is determined by minimizing electron group repulsions.
• Hierarchy of Electron Group Repulsions:
  • lone pair–lone pair > lone pair–bond pair > bond pair–bond pair
• Application:
  • VSEPR is instrumental in predicting geometries for compounds formed from non-metals.

3. Key Concepts of VSEPR

• Electron Group Geometry:
  • The 3D arrangement of electron groups in a molecule.
• Molecular Shape or Geometry:
  • The 3D arrangement of atoms in a molecule.
• Classification:
  • To classify molecular shape, assign each Lewis structure an AXₘEₙ designation, where:
  • A = central atom
  • X = surrounding bonded atom
  • E = non-bonding valence electron group
  • m, n = integers representing quantity
• Ideal Bond Angles:
  • Bond angles indicated are ideal (based solely on geometry).
  • Actual bond angles may vary due to influencing factors from surrounding electron groups.

4. VSEPR Notation and Molecular Shapes

5. Predicting Molecular Geometries

• Procedure Using VSEPR Theory
  • Draw the Lewis structure from the molecular formula.
  • Determine the total number of electron groups (both bonding and non-bonding) surrounding the central atom. Note: each multiple bond counts as a single electron group.
  • Identify the type of each electron group: bonding electron groups or lone pairs.
  • Establish the electron group arrangement around the central atom (linear, trigonal, tetrahedral, etc.).
  • Use established information and the table from prior slides to diagram and name the molecular shape or geometry.
  • For more complex molecules, assess the geometry of each central atom individually.

6. Example: Predicting the Molecular Shape of POCl₃

• Atom arrangement: Cl is typically a terminal atom, O is usually terminal, and P is the atom with the lowest electronegativity (EN).
• Valence Electrons Computation:
  • P: 6 + O: 5 + (3 × Cl: 7) = 32 electrons total.

7. Formal Charges Calculation for Molecular Structure

• For atoms:
  • P: $5 – ½(8) = +1$
  • Cl: $7 – [½(2) + 6] = 0$
  • O: $6 – [½(2) + 6] = -1$

8. Example: Nitrous Oxide, N₂O

• Application: Used as an anesthetic in dentistry.
• Prediction: Shape of N₂O molecule depends on its electron configuration.

9. Molecular Shapes and Dipole Moments (Section 10-5)

• Dipole Moment (µ):
  • Indicates the extent of charge displacement in polar bonds.
• Polar Covalent Bonds:
  • Exhibit bond dipole characteristics.
• Molecular Polarity:
  • Refers to molecules with a net charge imbalance that results in net dipole moment.

10. Determining Molecular Polarity in Multi-Atom Molecules

• For molecules with more than two atoms:
  • Both shape and bond polarity play roles in molecular polarity.
• CO₂ Example:
  • Two C=O bonds are polar; however, their linear arrangement leads to counterbalanced dipole moments, resulting in no net dipole moment.
• H₂O Example:
  • V-shaped molecular structure results in net dipole moment due to non-counterbalanced bond polarities.

11. Physicochemical Implication of Molecular Polarity

• Impact on Physical Properties:
  • Molecular polarity influences physical properties significantly, evidenced by measured dipole moments of selected compounds.

12. Bonding Theories Overview

• Requirements for a Bonding Theory:
  • Should explain observed properties like bond length, bond angles, bond dissociation energies, etc.
• Valence Bond Theory (VB):
  • Describes covalent bond formation via atomic orbital overlap.
  • Components of VB Method:
  • Valence electron configuration (e.g. Lewis structure).
  • Geometry predictions through VSEPR.
  • Molecular orbitals for lone and bonding electrons.

13. Key Principles of Valence Bond Theory

• Covalent Bond Formation:
  • An overlap of orbitals creates a bond occupied by a pair of electrons (highest probability located between nuclei).
• Themes Derived from Principal:
  • Opposing spins: electrons must have opposite spins (Pauli Exclusion Principle).
  • Maximum overlap: Strength of the bond correlates with the overlap of bonding orbitals.
  • Hybridization: The concept of mixing atomic orbitals to form new hybrid orbitals.

14. Hybrid Orbitals and Geometric Orientation

• Types of Hybrid Orbitals:
  • sp: linear
  • sp²: trigonal planar
  • sp³: tetrahedral
  • sp³d: trigonal bipyramidal
  • sp³d²: octahedral

15. Multiple Covalent Bonds

• Orbital Overlap Types:
  • End-to-end overlap results in a sigma (σ) bond.
  • Side-to-side overlap results in a pi (π) bond.
• Each multiple bond contains one σ and one or more π bonds.

16. Molecular Orbital Theory: Electron Delocalization

• Limitations of VB Model:
  • Assumes electrons are localized without resonance.
• Molecular Orbital Theory (MO):
  • Considers a molecule as a collection of nuclei with delocalized electron orbitals (LCAO method employed).
• Linear Combination of Atomic Orbitals (LCAO):
  • Each molecular orbital is formed from linear combinations of atomic orbitals.

17. Bonding and Antibonding Molecular Orbitals

• Description:
  • Bonding molecular orbitals (MOs) possess electron density between nuclei. Antibonding MOs do not.
• Electron Filling Order:
  • MOs filled following Aufbau principle, Hund's rule, and Pauli Exclusion Principle.

18. Bond Order and Its Implications

• Calculation:
  • Bond order = (number of electrons in bonding MO) - (number of electrons in antibonding MO).
• Observations:
  • Increased bond order correlates with increased bond energy and decreased bond length.

19. Example: Nitryl Fluoride

• Analysis based on mass percent composition and empirical formulas for predicting molecular shape and polarity.