Exam Study Notes

Ionic Compounds

• Forming ionic compounds requires energy to vaporize/ionize metals and atomize nonmetals.
• Electron affinity (EA) provides some energy back.
• Electrostatic attraction between oppositely charged ions in the crystal is the main driving force.
• Lattice energy is the heat emitted when gaseous ions coalesce to form a solid.
• Hess’s law is used to find lattice energy.
• E = \frac{q1q2}{4 \pi \epsilon_0 r}
• Estimating Lattice Energies:
  • Compare charges first: largest |q1q2| gives largest (negative) LE.
  • If q1q2 are the same, smaller radius ions give larger LE.

Covalent Bonding

• Covalent bonding involves sharing of electrons.
• A build-up of electron density between nuclei holds atoms together.
• Lewis Structure:
  • Represents bonding between atoms using Lewis symbols.
  • Does not indicate molecule shape.
• Octet Rule:
  • Atoms gain, lose, or share electrons to achieve eight valence electrons.
  • Hydrogen only needs 2 electrons (“Doublet” Rule).

Lewis Structures

• Determine the central atom (typically less electronegative).
• H is always a terminal atom.

Guidelines for Writing Lewis Structures

1. Add up valence electrons.
2. Create a skeleton structure with lines for bonding pairs.
3. Satisfy the Octet Rule for terminal atoms first.
4. Place remaining electrons around the central atom.
5. Move lone pairs to form multiple bonds if needed.

• HONC (for neutral organic molecules):
  • H: 1 bond
  • O: 2 bonds, 2 lone pairs
  • N: 3 bonds, 1 lone pair
  • C: 4 bonds

Electronegativity (EN)

• Pauling Scale indicates an atom's relative attraction for electrons.
• Trends: increases from lower left to upper right (excluding noble gases).
• Fluorine has the highest EN (4.0).
• Nonpolar covalent bonds: ΔEN ≈ 0 – 0.4
• Polar covalent bonds: ΔEN ≈ 0.4 – 2.0
• Ionic bonds: ΔEN > 2.0
• Electrostatic potential maps use colors to represent electron density.

Dipole Moments

• Represented by μ, measured in Debyes (D).
• μ = 4.77 \frac{D}{Å} × (bond \space length \space in \space Å)
• % ionic character = \frac{measured \space dipole \space moment}{calculated \space dipole \space moment} × 100%

Formal Charge

• Formal Charge is the charge on an atom if all shared electrons are equally shared.
• Formula: Formal Charge = (valence electrons) - (lone pair electrons) - (1/2 bonding electrons).
• Sum of formal charges is zero for a molecule, and equals the charge of an ion.
• Molecules obeying HONC have zero formal charges.
• Predicting Preferred Structure:
  • Smallest Formal Charges are most stable.
  • Negative Formal Charge should be on the most electronegative atom.