AP Chemistry Unit 2 ReviewCompound Structure and Properties

AP Chemistry Unit 2 is about how the structure of a compound, meaning how its atoms or ions are arranged and bonded, explains the properties you can actually measure. The single biggest idea is that electronegativity differences predict bond type, and Lewis structures plus VSEPR theory predict molecular shape and polarity, which together explain why substances behave the way they do. This unit is worth 7-9% of the AP exam, and its drawing-and-predicting skills (Lewis diagrams, geometry, polarity) show up constantly in later units too.

What this unit covers

Bond types and what electronegativity tells you

• Electronegativity increases left to right across a period and decreases down a group. You can explain both trends with Coulomb's law and the shell model. More protons pulling on the same shell means a stronger pull on electrons; more shells means valence electrons sit farther from the nucleus.
• The electronegativity difference between two atoms predicts the bond. Similar electronegativities give a nonpolar covalent bond (C-H counts as effectively nonpolar even though carbon is slightly more electronegative). A moderate difference gives a polar covalent bond with partial charges. A large difference, typically metal plus nonmetal, gives an ionic bond.
• Don't lean only on "metal + nonmetal = ionic." The exam expects you to use electronegativity values and properties as evidence, not just position on the periodic table.

Potential energy and bond strength

• A potential energy versus internuclear distance graph is the picture of a covalent bond. The bottom of the well marks the equilibrium bond length (where attraction and repulsion balance), and the depth of the well is the bond energy (how much energy it takes to pull the atoms apart).
• Bond order matters. A triple bond is shorter and stronger than a double bond, which is shorter and stronger than a single bond between the same atoms. Smaller atomic cores also mean shorter, stronger bonds.
• For ionic compounds, Coulomb's law does the same job. Higher charges and smaller ions mean stronger attractions and higher lattice energy. MgO (2+ and 2-) is held together much more tightly than NaCl (1+ and 1-).

Solids at the particle level: ionic crystals, metals, and alloys

• Ionic solids are a repeating 3-D lattice of cations and anions arranged to maximize attraction and minimize repulsion. That rigid lattice explains why ionic compounds are brittle, have high melting points, and only conduct electricity when melted or dissolved (the ions have to be free to move). You don't need to memorize specific crystal structures.
• Metallic bonding is a "sea of electrons" model. Positive metal ions sit in an array surrounded by delocalized valence electrons. Mobile electrons explain conductivity; ions that can slide past each other without shattering the structure explain malleability.
• Alloys come in two flavors. Interstitial alloys form when small atoms fill the gaps between larger ones (carbon in iron makes steel). Substitutional alloys form when atoms of similar size swap places in the lattice (copper and zinc make brass).

Lewis diagrams, resonance, and formal charge

• Lewis diagrams follow a set procedure. Count total valence electrons, connect atoms with bonds, distribute remaining electrons as lone pairs to satisfy octets, and form multiple bonds if the central atom comes up short.
• When more than one equivalent Lewis structure exists (like the three for NO3-), the real molecule is a resonance hybrid. The bonds are identical in length and strength, somewhere between single and double. No single structure is "the" molecule.
• When valid structures are NOT equivalent, formal charge breaks the tie. Formal charge equals valence electrons minus lone-pair electrons minus bonds. The best structure minimizes formal charges and puts any negative formal charge on the more electronegative atom.

VSEPR, hybridization, and polarity

• VSEPR theory is just Coulombic repulsion applied to electron pairs. Electron domains around a central atom spread out as far as possible, and the resulting geometry follows from how many bonding and lone pairs there are.
• Know the full geometry set: linear, trigonal planar, tetrahedral, trigonal pyramidal, bent, trigonal bipyramidal, seesaw, T-shaped, octahedral, and square pyramidal. Lone pairs repel more strongly than bonding pairs, which squeezes bond angles (water is 104.5 degrees, not 109.5).
• Hybridization follows from the geometry, not the other way around. Two electron domains means sp, three means sp2, four means sp3. Sigma bonds come from head-on orbital overlap; pi bonds come from side-by-side overlap in double and triple bonds.
• Molecular polarity needs two checks. A molecule is polar only if it has polar bonds AND a shape where the bond dipoles don't cancel. CO2 has two polar bonds but is nonpolar because they point in opposite directions. H2O is bent, so its dipoles add up.

Unit 2, Compound Structure and Properties at a glance

Why Unit 2, Compound Structure and Properties matters in AP Chem

This unit is the bridge from "what is an atom" to "why does this substance behave this way." AP Chemistry is built around explaining macroscopic properties with particle-level structure, and Unit 2 hands you the entire toolkit for doing that with compounds.

• It puts the course's structure-determines-properties theme front and center. Melting point, conductivity, and solubility arguments all start with the bonding and shape ideas built here.
• Coulomb's law shows up as the explanation engine behind almost everything: electronegativity trends, lattice energy, bond strength, and VSEPR repulsion are all the same physics in different costumes.
• Lewis diagrams and VSEPR are not just Unit 2 skills. They are prerequisites for predicting intermolecular forces, acid strength, and reaction behavior later, and free-response questions assume you can draw and interpret them quickly.

How this unit connects across the course

• Electronegativity trends and Coulomb's law come straight out of atomic structure and periodicity (Unit 1). If the periodic trends feel shaky, review them first, because Unit 2 uses them as evidence in every bonding argument.
• Molecular polarity from VSEPR is the setup for intermolecular forces, which control boiling points, solubility, and chromatography (Unit 3). You can't rank IMFs without knowing whether a molecule is polar.
• Bond energy from the potential energy curve returns when you estimate reaction enthalpies from bonds broken and bonds formed (Unit 6). Breaking bonds costs energy; forming bonds releases it.
• Lewis structures and electron-pair thinking pay off in acid-base chemistry (Unit 8), where structure explains why some acids are stronger than others, and in reaction types and net ionic equations (Unit 4).

Key equations and processes

• Coulomb's law, E ∝ (q1 × q2)/r. Use it to compare lattice energies, explain electronegativity trends, and justify why smaller, more highly charged ions attract more strongly.
• Formal charge = (valence electrons) − (lone-pair electrons) − (number of bonds). Use it to choose the best Lewis structure when nonequivalent options exist.
• The Lewis diagram procedure: count total valence electrons (adjust for ion charge), bond the atoms, fill octets with lone pairs, then convert lone pairs to multiple bonds if needed.
• Reading a potential energy curve: the minimum gives equilibrium bond length, and the depth of the well gives bond energy. Compare wells to compare bond strengths.
• VSEPR domain counting: count electron domains on the central atom (lone pairs and bonds each count once, and a double or triple bond counts as one domain), then assign the geometry and adjust angles for lone-pair repulsion.
• Hybridization from domains: 2 domains = sp, 3 = sp2, 4 = sp3. Use it to describe sigma and pi bonding in a structure.

Unit 2, Compound Structure and Properties on the AP exam

Unit 2 is worth 7-9% of the exam, but its skills are baked into far more questions than that number suggests. Multiple-choice questions ask you to pick the correct particulate diagram of an ionic solid or alloy, identify bond type from electronegativity data, interpret potential energy curves, and match a formula to its geometry and polarity. Free-response questions routinely tell you to draw a complete Lewis diagram (including all lone pairs), then use it to justify a geometry, a bond angle, a hybridization, or a polarity claim in a follow-up part. The pattern to practice is draw, then explain. A correct shape with no Coulombic or VSEPR reasoning attached usually earns only partial credit, so get comfortable writing one or two sentences that connect the structure to the property. Resonance questions often ask you to explain why all the bonds in an ion like carbonate are the same length, and formal charge questions ask you to defend a choice between two valid diagrams.

Essential questions

• How does the arrangement of atoms and electrons in a compound determine the properties we can measure in the lab?
• Why do some atoms transfer electrons while others share them, and how can you predict which will happen?
• How can a 2-D Lewis diagram tell you the 3-D shape and polarity of a molecule?
• Why do ionic solids, metals, and molecular substances behave so differently even though they're all "just atoms"?

Key terms to know

• Electronegativity: A measure of how strongly an atom attracts shared electrons in a bond.
• Polar covalent bond: A bond where electrons are shared unequally, creating partial positive and negative charges on the atoms.
• Bond order: The number of shared electron pairs between two atoms; higher bond order means a shorter, stronger bond.
• Bond energy: The energy required to separate two bonded atoms, shown as the depth of the potential energy well.
• Lattice energy: The energy associated with separating the ions of an ionic crystal, governed by Coulomb's law.
• Sea of electrons: The model of metallic bonding where delocalized valence electrons move freely among an array of metal cations.
• Interstitial alloy: An alloy where small atoms fill the spaces between larger atoms, like carbon in iron to make steel.
• Substitutional alloy: An alloy where atoms of similar size replace each other in the lattice, like brass.
• Resonance: The situation where a molecule's true structure is a hybrid of two or more equivalent Lewis structures.
• Formal charge: A bookkeeping value (valence electrons minus lone-pair electrons minus bonds) used to pick the best Lewis structure.
• VSEPR theory: The idea that electron pairs around a central atom repel each other and spread out, which sets molecular geometry.
• Hybridization: The mixing of atomic orbitals (sp, sp2, sp3) that matches the number of electron domains around an atom.
• Sigma and pi bonds: Sigma bonds form from head-on orbital overlap; pi bonds form from side-by-side overlap in double and triple bonds.
• Dipole moment: The overall separation of charge in a molecule, present only when bond dipoles don't cancel.

Common mix-ups

• Electron geometry vs. molecular geometry. Lone pairs count when finding the electron arrangement but are invisible in the molecular shape name. NH3 has tetrahedral electron geometry but trigonal pyramidal molecular geometry.
• Polar bonds don't guarantee a polar molecule. CO2 and CCl4 both have polar bonds, but their symmetric shapes cancel the dipoles. Always check the geometry before calling a molecule polar.
• Resonance structures are not molecules flipping back and forth. The real species is one hybrid with identical, intermediate bonds, which is why all three N-O bonds in nitrate are the same length.
• Bond energy and lattice energy answer different questions. Bond energy is about pulling apart two covalently bonded atoms; lattice energy is about separating all the ions in an ionic crystal. Both get stronger with higher charge and shorter distance, but they apply to different bonding types.