AP Chemistry Unit 2 Cheatsheet: Compound Structure and Properties

Electronegativity trends:

• Increases left → right across period
• Decreases down a group
• Explained by atomic size and nuclear charge (Coulomb's law)

Bond classification by electronegativity difference:

• Nonpolar covalent: Similar EN (e.g., C–H)
• Polar covalent: Unequal EN → partial charges (δ+ and δ−)
• Ionic: Metal + nonmetal; large EN difference

Key principle: Ionic–covalent bonding is a continuum, not distinct categories. Greater EN difference = stronger bond dipole in single bonds. Examine compound properties to confirm bond type.

PE vs. internuclear distance graph shows:

• Equilibrium bond length: Distance at minimum PE
• Bond energy: Energy required to separate atoms (break bond)

Factors affecting bond strength:

• Bond order: Triple > double > single (shorter, higher energy)
• Atomic size: Smaller atoms = shorter bonds, higher energy
• Coulomb's law (ionic): Strength ∝ (charge₁ × charge₂) / distance²

Ionic interaction strength: Larger charges and smaller ions → stronger attractions.

Metallic bonding model: Array of cations surrounded by delocalized "sea" of valence electrons (not bound to individual atoms).

Interstitial alloys:

• Atoms of significantly different radii
• Smaller atoms fill spaces between larger atoms
• Example: carbon in iron (steel)

Substitutional alloys:

• Atoms of comparable radius
• One atom replaces another in lattice
• Example: zinc in copper (brass)