Lecture Notes: Sections 13.2 and 13.3

Review of Ionic and Covalent Bonds

Ionic Bonds
- Formed by a cation and an anion.
- Often involves a metal and a non-metal.
- Can involve polyatomic ions.
Covalent Bonds
- Formed between two non-metals.
- Involves sharing of electrons.
- Examples include H-H, Cl-Cl, O-O.

Hydrogen Atoms Interaction
- Two hydrogen atoms (A and B): start infinitely apart.
- Energy initially set at zero.
- As atoms are brought closer:
  - Electron from A sees proton in B and vice versa.
  - Favorable interaction lowers system energy.
  - Optimal bond length minimizes proton-proton repulsion.
- Takeaway: Sharing electrons lowers system energy, forming a stable covalent bond.

Homonuclear Molecules
- Same atoms form non-polar molecules (e.g., H-H, Cl-Cl).
Heteronuclear Molecules
- Different atoms can lead to polar bonds.
- Example: HCl
  - When placed between charged plates, aligns due to dipole nature.
  - Electrons not shared equally; chlorine pulls electrons more (greedier).
  - Results in a dipole arrow pointing to the more electronegative atom (Cl).
  - Partial negative charge (δ-) on Cl, partial positive charge (δ+) on H.
Polar vs. Non-polar
- Polar: Unequal sharing, creates dipole.
- Non-polar: Equal sharing, no dipole.

Definition: Ability of an atom to attract electrons in a compound.
- Different from electron affinity (adding an electron to an atom).
Trends in Electronegativity
- Fluorine: Most electronegative.
- Decreases down a group and from right to left.
Delta Electronegativity
- Calculate by subtracting values of two atoms.
- Greater difference indicates ionic character.
Bond Type Classifications
- Ionic: ΔEN > 1.8
- Polar Covalent: 0.4 < ΔEN ≤ 1.8
- Mostly Covalent: ΔEN < 0.4

HCl: Polar covalent bond (ΔEN = 0.9)
NaCl: Ionic bond (ΔEN = 2.1)
LiH: Polar covalent bond (ΔEN = 1.1)
HF: Polar covalent bond (ΔEN = 1.9)
RbO: Most ionic (ΔEN = 2.7)