AP Chemistry · Handsworth Secondary 2026–27
Unit 2 · Study Guide
Dr. Ras Mulinta
Compound Structure & Properties
Exam-focused review
A one-page map of what the Unit 2 test (and the AP exam) expects. This is a checklist, not a re-teach, if a line doesn't click, go back to that section of the notes package. A periodic table is provided; bond data are given when needed.
Must be able to do
- 2.1 Use electronegativity & periodic position to classify a bond as ionic, polar covalent, nonpolar covalent, or metallic, and tie each to physical properties.
- 2.2 Read a potential-energy vs. internuclear-distance curve: minimum = bond length, well depth = bond energy. Higher bond order → shorter, stronger.
- 2.3 Model an ionic solid as an alternating 3-D lattice; explain hardness, high melting point, brittleness, and conductivity with Coulomb's law.
- 2.4 Use the "sea of electrons" to explain metallic properties; distinguish interstitial vs. substitutional alloys by atomic radius.
- 2.5 Draw correct Lewis diagrams (right valence-electron count, octets, multiple bonds) for molecules and polyatomic ions.
- 2.6 Compute FC = V − L − ½B to pick the best structure; draw resonance structures and give fractional bond orders.
- 2.7 From a Lewis diagram: electron & molecular geometry, bond angles, polarity (vector cancellation), hybridization (sp/sp²/sp³), and σ vs. π counts.
The big idea that ties it together
Structure dictates properties. Electronegativity sets the bond type; the Lewis diagram + VSEPR set the shape; the shape sets the polarity, the hybridization, and the bond order, and those, through F ∝ q₁q₂/r², set what you measure: melting point, conductivity, bond length, bond strength. Draw it right and the macroscopic behaviour falls out.
Don't waste time on (excluded by the CED)
Specific named crystal structures (rock-salt vs. CsCl lattices) · the derivation/depiction of hybrid orbitals · d-orbital hybridization (for 5–6 domains give only the shape) · molecular-orbital diagrams & bonding/antibonding filling · detailed radical (odd-electron) bookkeeping. Know the models and the patterns; skip the orbital math.
Quick self-check (answers in your head, then verify)
- Classify the bonding in KBr, F₂, and Al: ionic, covalent, or metallic?
- Give the electron geometry, molecular geometry, and bond angle of NH₃.
- What is the bond order of each N–O bond in the nitrate ion, NO₃⁻?
- Compute the formal charge on the central O in ozone drawn as O=O–O (double bond left, single bond right, one lone pair on the centre).
- How many σ and how many π bonds are in hydrogen cyanide, H–C≡N?
- Is BF₃ polar or nonpolar? Justify with its shape.
- Rank C–C, C=C, and C≡C by bond length (longest first).
Check yourself: 1) KBr ionic (metal + nonmetal), F₂ nonpolar covalent (ΔEN = 0), Al metallic. 2) electron geometry tetrahedral, molecular geometry trigonal pyramidal, angle ≈ 107°. 3) bond order = 4 bonds ÷ 3 bonds = 1.33 (one double shared over three equivalent N–O bonds). 4) FC = 6 − 2 − ½(6) = +1. 5) 2 σ and 2 π (C–H = 1 σ; C≡N = 1 σ + 2 π). 6) nonpolar, trigonal planar with three identical B–F bonds, so the bond dipoles cancel. 7) C–C > C=C > C≡C (higher bond order = shorter bond).