AP Chemistry Unit 2 Review: Compound Structure and Properties

Every Unit 2 topic, what's actually tested, and what the exam explicitly excludes.

Unit 2 takes the atomic structure from Unit 1 and asks how atoms combine into compounds -- and how that combination determines a substance's properties. It's 7–9% of the multiple-choice section across ~12–13 class periods, but Lewis structures and VSEPR geometry are tools you'll reuse constantly in Unit 3 (intermolecular forces) and beyond. Here's what's actually in it, topic by topic, verified against the current CED -- including four exclusion statements that cut real content most courses still teach.

Unit 2 at a Glance

Unit 2 (Compound Structure and Properties) covers 7 topics. See the full AP Chemistry units breakdown for how it fits alongside the other 8 units, and the Course and Exam Description for the full framework.

Types of Chemical Bonds

Electronegativity increases left-to-right across a period and decreases down a group. That trend determines the type of bond formed:

Nonpolar covalent bonds form between atoms of similar electronegativity -- C–H bonds count as effectively nonpolar even though carbon is slightly more electronegative than hydrogen.
Polar covalent bonds form between atoms of unequal electronegativity. The more electronegative atom develops a partial negative charge, and larger electronegativity differences produce larger bond dipoles.
Ionic bonds generally form between a metal and a nonmetal.

Common misconception: electronegativity difference alone doesn't determine ionic versus covalent bonding -- the ionic/covalent distinction is a continuum, not a sharp line. The CED is explicit that examining a compound's actual properties is the best way to characterize its bonding, not just running a Δ-electronegativity number.

Intramolecular Force and Potential Energy

A graph of potential energy versus internuclear distance describes how two atoms interact as they approach each other -- the minimum of the curve is the equilibrium bond length, and the well depth is the bond energy.

Bond order affects both: higher-order bonds (double, triple) are shorter and have greater bond energy than single bonds. For ionic interactions specifically, Coulomb's law applies directly -- larger ionic charges and smaller ionic radii both produce stronger (more negative) interaction energies.

Structure of Ionic Solids

Cations and anions arrange in a systematic, periodic 3-D array that maximizes attractive forces between opposite charges while minimizing repulsive forces between like charges.

Not tested: knowledge of specific crystal structures (rock-salt, cesium-chloride, and similar named lattice geometries) is explicitly excluded. You only need the general principle above, not memorized unit cells.

Structure of Metals and Alloys

Metallic bonding is a lattice of positive metal ions surrounded by delocalized valence electrons -- a "sea of electrons." Two alloy types follow from this:

Interstitial alloys form between atoms of significantly different radii, where the smaller atoms fill gaps between the larger ones (e.g., carbon in steel).
Substitutional alloys form between atoms of comparable radius, where one atom directly replaces another in the lattice (e.g., zinc substituting for copper in brass).

Lewis Diagrams

Lewis diagrams are constructed following an established set of principles: total valence electrons (adjusted for ion charge), a central atom bonded to surrounding atoms, and remaining electrons distributed as lone pairs to satisfy the octet rule wherever possible.

Worked example: Draw the Lewis structure for the nitrate ion, NO₃⁻. Total valence electrons: N (5) + 3 × O (6) + 1 (for the negative charge) = 24 electrons. Place N centrally bonded to three O atoms (3 single bonds = 6 electrons), distribute the remaining 18 electrons as lone pairs on the oxygens, then convert one lone pair into a second N=O bond to give nitrogen a complete octet. The result is one N=O double bond and two N–O single bonds -- and because those three positions are equivalent, resonance applies (see below).

Resonance and Formal Charge

When more than one equivalent Lewis structure can be drawn, resonance must be included -- the real molecule is a blend of the contributing structures, not any single one of them. When structures aren't equivalent, formal charge (FC = valence electrons − nonbonding electrons − bonding electrons) and the octet rule together determine which structure best predicts molecular behavior.

Worked example: In one valid Lewis structure for SO₂, sulfur carries one lone pair, one S=O double bond, and one S–O single bond. Using formal charge = valence electrons − nonbonding electrons − ½(bonding electrons): sulfur = 6 − 2 − ½(6) = +1; the singly-bonded oxygen (3 lone pairs) = 6 − 6 − ½(2) = −1; the doubly-bonded oxygen (2 lone pairs) = 6 − 4 − ½(4) = 0. The charges sum to zero, as they must for a neutral molecule. Because an equivalent structure exists with the double bond on the other oxygen, SO₂ is best represented as a resonance hybrid of both contributors rather than one fixed structure.

Model limitation: the CED explicitly notes that the Lewis structure model has limitations, particularly for species with an odd number of valence electrons -- don't expect every molecule to fit cleanly into the octet-rule framework.

VSEPR and Hybridization

VSEPR theory uses Coulombic repulsion between electron pairs to predict molecular geometry. Combined with Lewis diagrams, it predicts molecular geometry, bond angles, relative bond energies and lengths, dipole moments, and orbital hybridization for covalently bonded molecules and polyatomic ions.

Worked example: Predict the molecular geometry of NH₃. Nitrogen has 5 valence electrons; three go into N–H bonds, leaving one lone pair. Four electron domains around N (3 bonding + 1 lone pair) means an electron-domain geometry of tetrahedral, but because one domain is a lone pair, the molecular geometry is trigonal pyramidal, with bond angles slightly less than 109.5° since lone-pair repulsion compresses the bonding angles.

Bond formation involves orbital overlap: sigma bonds (stronger, from direct end-to-end overlap) and pi bonds (weaker, from side-to-side overlap in multiple bonds). A pi bond also locks rotation around that bond, which is why double bonds can produce geometric (cis/trans) isomers.

Not tested (three exclusions in this topic alone): (1) the derivation and depiction of hybrid orbitals is excluded -- you need the sp/sp²/sp³ labels and angles, not orbital-mixing diagrams; (2) hybridization involving d orbitals is excluded -- for more than four electron domains, you're only responsible for the resulting molecular shape; (3) molecular orbital theory is excluded entirely -- no MO diagrams, no filling of molecular orbitals, no bonding/nonbonding/antibonding distinctions.

Common Mistakes in Unit 2

Drawing hybrid-orbital mixing diagrams to "prove" hybridization. That derivation is explicitly excluded -- state the hybridization label and angle, and move on.
Reaching for molecular orbital theory language (bonding/antibonding) to explain bond strength. MO theory isn't assessed; explain bond order and overlap using Lewis/VSEPR language instead.
Treating electronegativity difference as the sole test for ionic vs. covalent bonding. The CED is explicit that it's a continuum best confirmed by a compound's actual properties.
Confusing electron-domain geometry with molecular geometry. A tetrahedral electron-domain arrangement with one lone pair gives a trigonal pyramidal molecular shape (like NH₃) -- the two labels aren't interchangeable.
Forgetting to adjust valence electron count for ion charge. Add one electron per negative charge, subtract one per positive charge, before drawing any Lewis structure for an ion.

How Unit 2 Connects to the Rest of the Course

The Lewis structures and VSEPR geometries built here are the starting point for predicting Unit 3's intermolecular forces -- you can't identify a dipole moment without a correct molecular geometry.
Formal charge reasoning resurfaces directly when comparing resonance structures for acids and bases in Unit 8.
The potential-energy-vs-distance curve from this unit is the same framework used for reaction energy profiles in Unit 5 (Kinetics).

Related Resources

Frequently Asked Questions

What topics are in AP Chemistry Unit 2?

Types of Chemical Bonds, Intramolecular Force and Potential Energy, Structure of Ionic Solids, Structure of Metals and Alloys, Lewis Diagrams, Resonance and Formal Charge, and VSEPR and Hybridization -- 7 topics in total.

How much is Unit 2 worth on the AP Chemistry exam?

Seven to nine percent of the multiple-choice section, but Lewis structures and VSEPR reasoning resurface constantly in later units, especially Unit 3 intermolecular forces.

Does AP Chemistry require drawing hybrid orbital diagrams?

No. The derivation and depiction of hybrid orbitals is explicitly excluded. You do need the sp, sp2, and sp3 nomenclature and their associated bond angles, but not orbital-mixing diagrams.

Does AP Chemistry test molecular orbital theory?

No. Molecular orbital diagrams, the filling of molecular orbitals, and the distinction between bonding, nonbonding, and antibonding orbitals are all explicitly excluded from the exam.

Do I need to know specific ionic crystal structures for AP Chemistry?

No. Knowledge of specific crystal structures (like rock-salt or cesium-chloride geometry) is explicitly not required. You only need to explain that cations and anions arrange in a repeating 3-D pattern that maximizes attraction and minimizes repulsion.

Sourced from College Board's official AP Chemistry Course and Exam Description, Effective Fall 2024. This page describes the document's real content and current exclusion statements; it is not a copy of it and is not affiliated with or endorsed by College Board.