In-Depth Notes on Simple Bonding Theories

Chemical Bonding Overview

  • Types of Chemical Bonds
    • Ionic Bonds
      • Formed by electrostatic attraction between oppositely charged ions.
      • Requires a significant difference in electronegativity (EN) (usually > 2.0).
    • Covalent Bonds
      • Formed by the sharing of electrons between atoms.
      • EN difference is small or nearly negligible (< 2.0).

Dipole Moments and Polar Molecules

  • Dipole Moment: A quantitative measure of charge separation in a molecule, denoted as ( \mu ).
    • Formula: ( \mu = Q \times r ) where ( Q ) is the charge and ( r ) is the distance.
    • Units: Debyes (D), where ( 1 D = 3.336 \times 10^{-30} ) coulomb meter.
    • Polarity Criteria:
      • If ( \mu = 0 ): Non-polar molecule.
      • If ( \mu > 0 ): Polar molecule.

Drawing Lewis Structures

  1. Skeletal Structure: Start with the atoms connected, placing the least electronegative atom in the center.
  2. Valence Electrons: Count total valence electrons. Adjust for charges (add one for negatives, subtract one for positives).
  3. Octet Rule: Complete octets for surrounding atoms (except Hydrogen), and place remaining electrons on the central atom.
  4. Bonds: If fewer than 8 electrons on the central atom, form double or triple bonds by moving electron pairs from surrounding atoms.

Formal Charge in Lewis Structures

  • Formal Charge: Difference between the number of valence electrons in a free atom and those assigned in a Lewis structure.
    • Preferred structures have no formal charges or smaller formal charges.
    • Negative charges should reside on more electronegative atoms when possible.
    • The sum of formal charges must equal the ion or molecule's charge.
  • Example of Formal Charge Calculation:
    • For Nitric acid, HNO3:
      • ( \text{Formal Charge}_{H} = 1 - 0 - 1 = 0 )
      • ( \text{Formal Charge}_{N} = 5 - 0 - 4 = +1 )
      • ( \text{Formal Charge}_{O} = 6 - 6 - 1 = -1 )

Resonance Structures

  • Resonance: More than one valid structure can be drawn for a molecule.
    • Necessary to depict each structure to describe electron distribution adequately.
    • Example: Nitrate ion ( NO3^- ) has three equivalent structures depicting nitrogen-oxygen bonds.

Invalid Electron-Dot Structures

  • Requires adherence to the octet rule; exceeding it results in invalid structures.
  • Example: Structures exceeding octets cannot be considered valid.

Expanded Octet and Electron-Deficient Molecules

  • Expanded Octet: Atoms from the 3rd period and below can accommodate more than 8 electrons due to the availability of d orbitals.
    • Common in elements like ClF3 and SF4.
  • Electron-Deficient Molecules: Be and B often have fewer than 8 electrons; thus, they are highly reactive.
    • For stabilization: Lewis structures prioritize octets for heavier elements first.