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Unit 2 · Compound Structure & Properties Unit Hub Flashcards Cheat Sheet Essentials Visual Review MC Practice SAQ Practice

AP Chemistry Unit 2 Visual Review

A topic-by-topic visual walkthrough of Compound Structure & Properties — bond types, ionic and metallic solids, Lewis diagrams, resonance, and VSEPR.

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TOPIC 2.1 Types of Chemical Bonds BONDING IS A SPECTRUM — SET BY ELECTRONEGATIVITY DIFFERENCE (ΔEN) nonpolar covalent (ΔEN ≈ 0) polar covalent ionic (large ΔEN) Covalent bonds Two nonmetals SHARE electron pairs. Equal sharing = nonpolar (H₂, O₂); unequal = polar (HCl). Forms molecules; low melting points; poor conductors. Ionic bonds A metal TRANSFERS electrons to a nonmetal → oppositely charged ions attract (NaCl, MgO). High melting points; brittle; conduct when molten/dissolved. Metallic bonds Metal atoms share a mobile "sea" of delocalized valence electrons (Cu, Fe, Al). Malleable, ductile, shiny, and excellent conductors. A compound's bond type is the single best predictor of its physical properties (melting point, conductivity, hardness). The electronegativity difference decides the bond: shared (covalent), transferred (ionic), or pooled (metallic). The Review Hub · AP Chemistry Unit 2 TOPIC 2.2 Intramolecular Force & Potential Energy bond length bond energy (depth) potential energy distance between nuclei → The bond forms at the energy minimum As two atoms approach, attraction lowers PE — until they get too close and repulsion spikes it. The bottom of the well is the stable bond: the bond length & bond energy. Stronger bonds = deeper, shorter Bond order tells the story: triple > double > single bond. More shared pairs → SHORTER length and STRONGER (higher) bond energy. Example: C≡C (837) > C=C (614) > C–C (347 kJ/mol). Coulomb again The depth of the well (bond strength) is Coulombic: greater nuclear charge and shorter distance → stronger attraction. Breaking a bond ABSORBS energy (endothermic); forming a bond RELEASES it (exothermic). A bond sits at the minimum of the PE curve — higher bond order means shorter and stronger. The Review Hub · AP Chemistry Unit 2 TOPIC 2.3 Structure of Ionic Solids IONIC LATTICE (NaCl) Na⁺ Cl⁻ alternating 3-D array A repeating lattice of alternating ions Ions pack into a regular 3-D crystal lattice so that every ion is surrounded by oppositely charged neighbors, maximizing attraction and minimizing repulsion. Properties explained by the lattice • High melting points — strong Coulombic forces throughout • Brittle — shifting the lattice aligns like charges → it cracks • Conduct only when molten or dissolved — ions must be free Solid ionic compounds do NOT conduct — the ions are locked in place. Higher ion charges and smaller ions give a larger lattice energy → higher melting point (MgO > NaCl). Ionic solids form a rigid lattice — explaining their high melting points, brittleness, and conductivity only when free. The Review Hub · AP Chemistry Unit 2 TOPIC 2.4 Structure of Metals & Alloys "SEA OF ELECTRONS" MODEL ++++ ++++ fixed cations in a sea of mobile e⁻ (teal) Delocalized electrons explain metal properties • Conductive — free electrons carry charge & heat • Malleable & ductile — cations slide without breaking bonds • Lustrous — mobile electrons reflect light Alloys — metal mixtures Substitutional — similar-sized atoms replace host atoms in the lattice (e.g., brass = copper + zinc). Interstitial — small atoms fit in the gaps (e.g., steel = iron + carbon), making the metal harder & stronger. Alloying disrupts the regular lattice, so alloys are usually harder and less malleable than pure metals. A sea of delocalized electrons makes metals conductive, malleable, and shiny; alloys tune those traits. The Review Hub · AP Chemistry Unit 2 TOPIC 2.5 Lewis Diagrams LEWIS STRUCTURE OF CO₂ O C O two double bonds · each O has 2 lone pairs total valence e⁻ = 4 + 2(6) = 16 How to draw a Lewis structure 1. Count total valence electrons (sum over all atoms; add for negative charge, subtract for positive) 2. Place the least electronegative atom in the center 3. Connect atoms with single bonds (2 e⁻ each) 4. Add lone pairs to complete octets on outer atoms 5. If the center lacks an octet, form double/triple bonds H only wants 2 electrons. Some atoms (B, expanded octets) are exceptions. The octet rule Most main-group atoms are most stable with 8 valence electrons (a full s + p shell), like a noble gas. A shared pair (bond) counts toward BOTH atoms' octets. Lone pairs count only for the atom they sit on. Lewis structures predict bonding, but VSEPR (Topic 2.7) predicts the actual 3-D shape. Lewis diagrams map valence electrons — count them, connect atoms, and complete every octet. The Review Hub · AP Chemistry Unit 2 TOPIC 2.6 Resonance & Formal Charge Resonance structures When more than one valid Lewis structure can be drawn (by moving double bonds), the real molecule is an AVERAGE — a "hybrid" — of all of them. O=N–O⁻ ↔ ⁻O–N=O both N–O bonds are actually identical Formal charge picks the best structure FC = valence e⁻ − (lone e⁻ + ½ bonding e⁻) The BEST Lewis structure has formal charges closest to zero, and any negative charge on the most electronegative atom. Lower (near-zero) formal charges = more stable structure. Worked example — formal charge on an atom Find the formal charge of the doubly-bonded O in the nitrate ion (with 2 lone pairs). Valence e⁻ of O = 6 · lone e⁻ = 4 · bonding e⁻ = 4 (a double bond) FC = 6 − (4 + ½·4) = 6 − 6 = 0 A singly-bonded O with 3 lone pairs: FC = 6 − (6 + 1) = −1. The sum of all formal charges = the ion's charge. Formal charge is a bookkeeping tool — it is not the same as oxidation number. Resonance = an average of valid structures; formal charges near zero pick the best one. The Review Hub · AP Chemistry Unit 2 TOPIC 2.7 VSEPR & Bond Hybridization Electron domains (bonds + lone pairs) repel and spread out as far as possible → the molecular shape 2 domains → Linear 180° · e.g., CO₂ hybridization: sp 3 → Trigonal planar 120° · e.g., BF₃ · sp² 4 → Tetrahedral 109.5° · e.g., CH₄ · sp³ Lone pairs bend the shape Lone pairs repel MORE than bonds, so they compress bond angles and change the molecular shape: • 4 domains, 1 lone pair → trigonal pyramidal (NH₃, ~107°) • 4 domains, 2 lone pairs → bent (H₂O, ~104.5°) Count electron domains first, THEN subtract lone pairs for the shape you actually see. Shape → polarity A molecule is POLAR if it has polar bonds AND an asymmetrical shape (dipoles don't cancel). • H₂O is bent & polar · NH₃ is pyramidal & polar • CO₂ is linear & NONpolar (dipoles cancel) Hybridization: count domains — 2 = sp, 3 = sp², 4 = sp³. Molecular polarity drives intermolecular forces (Unit 3). Electron domains repel to set the shape and bond angles — which then determine polarity. The Review Hub · AP Chemistry Unit 2
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How to use the visual review

Spend 30 seconds per slide before clicking next. Look at the diagram, then ask yourself: "Could I draw this from memory and explain it?"

Use the fullscreen button () on desktop for the best experience. Use arrow keys to navigate. Tap "Show all slides" to jump around.

This is great for review the night before the exam — fast, visual, and covers everything you need to remember about Unit 2's bonding and molecular structure content.