Chemical Bonding Notes

Formal Charge

• Formal Charge = # of valence electrons - # of assigned electrons
• Assigned electrons:
  • Lone pair electrons are assigned to the atom.
  • Bonding pairs are equally shared (one electron to each atom).
• Sum of formal charges:
  • Zero for a molecule.
  • Equal to the charge on an ion.
• Molecules obeying HONC have all zero formal charges.

Predicting Preferred Structure

• Smallest Formal Charges is most stable
• If equal, negative Formal Charge on the most electronegative atom is more stable.

Resonance

• Two or more valid electron dot structures differing only in electron arrangement.
• No single Lewis structure correctly represents bonding.
• True structure is an average of resonance forms.
• Molecules with resonance have added stability.
• Equivalent structures contribute equally; non-equivalent structures contribute based on stability.

Exceptions to the Octet Rule

• Less than an octet:
  • Be (usually 4 electrons).
  • B (frequently 6 electrons).
  • Avoid double bonds between Be/B and F/Cl.
• Odd number of electrons: Reactive free radicals (e.g., NO, NO2, ClO2).
• Expanded octet:
  • Elements in the second period (B, C, N, O, F) never exceed an octet.
  • Elements in the third period and below can exceed an octet due to available d orbitals.

Bond Energies

• Bonds of a given type have approximately the same energy (gas phase only).
• Breaking bonds is endothermic; forming bonds is exothermic.
• \Delta H = \Sigma (bonds broken) - \Sigma (bonds formed)

Molecular Shapes and VSEPR Theory

• Molecular shape is determined by electron groups around each atom.
• VSEPR: Valence Shell Electron Pair Repulsion Theory
  • Draw Lewis structure.
  • Count electron groups around the central atom (double/triple bonds count as one group).
  • Count lone pairs.
  • Use number and type of electron groups to predict shape.