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Unit 9 · Thermodynamics & Electrochemistry Unit Hub Flashcards Cheat Sheet Essentials Visual Review MC Practice SAQ Practice

AP Chemistry Unit 9 Visual Review

A topic-by-topic visual walkthrough of Thermodynamics & Electrochemistry — entropy, Gibbs free energy, galvanic and electrolytic cells.

← Back to Unit 9 hub
TOPIC 9.1 Introduction to Entropy Entropy (S) = dispersal of energy Entropy measures how spread out matter and energy are — the number of ways a system can be arranged. More microstates ⇒ higher entropy. Entropy INCREASES when… • solid → liquid → gas • a solid/liquid dissolves • moles of GAS increase in a reaction • temperature rises Low S → high S ordered (low S) → dispersed (high S) Gases spread to fill their container — many more arrangements available. Entropy = energy & matter dispersal — it rises with more gas, higher T, and dissolving. The Review Hub · AP Chemistry Unit 9 TOPIC 9.2 Absolute Entropy and Entropy Change ΔS°_rxn = ΣS°(products) − ΣS°(reactants) Using standard entropies S° values (J/mol·K) are tabulated for each substance. Weight each by its coefficient, then subtract. Predict the sign first: if gas moles rise, ΔS is usually positive. Absolute entropy & the 3rd law A perfect crystal at 0 K has S = 0 (only one microstate). So every substance has a positive absolute S° above 0 K. Larger, more complex molecules → higher S°. ΔS° = ΣS°(products) − ΣS°(reactants); predict the sign from changes in gas moles. The Review Hub · AP Chemistry Unit 9 TOPIC 9.3 Gibbs Free Energy & Thermodynamic Favorability ΔG° = ΔH° − TΔS° WHEN IS ΔG NEGATIVE (FAVORABLE)? ΔH −, ΔS + Favorable at ALL temperatures ΔH +, ΔS − Never favorable (nonspontaneous at all T) ΔH −, ΔS − Favorable at LOW temperature only ΔH +, ΔS + Favorable at HIGH temperature only A process is thermodynamically favorable when ΔG < 0 — temperature can flip the sign. The Review Hub · AP Chemistry Unit 9 TOPIC 9.4 Thermodynamic and Kinetic Control Favorable ≠ fast ΔG < 0 tells you a reaction CAN happen, not how fast it will. Rate is governed by activation energy (kinetics). Thermodynamics = direction; kinetics = speed. Kinetically stable A large activation energy can make a favorable reaction immeasurably slow. Example: diamond → graphite has ΔG < 0 but essentially never occurs. A catalyst changes rate, not favorability A catalyst lowers Eₐ so a favorable reaction proceeds faster — but it does NOT change ΔG, ΔH, or K. It cannot make an unfavorable reaction (ΔG > 0) become favorable; it only speeds up reaching equilibrium. Same start and end points → same thermodynamics. ⏱Thermodynamics says whether; kinetics says how fast — catalysts change only the speed. The Review Hub · AP Chemistry Unit 9 TOPIC 9.5 Free Energy and Equilibrium ΔG° = −RT ln K ΔG° < 0 K > 1 Products favored at equilibrium. Reaction lies to the right (more products). ΔG° = 0 K = 1 Comparable amounts of both sides. Neither side strongly favored. ΔG° > 0 K < 1 Reactants favored at equilibrium. Reaction lies to the left (more reactants). ΔG° = −RT ln K — negative ΔG° means K > 1 and products are favored. The Review Hub · AP Chemistry Unit 9 TOPIC 9.6 Free Energy of Dissolution Dissolving is a balance of ΔH and ΔS Whether a salt dissolves is decided by ΔG = ΔH − TΔS for the dissolution process. Both terms can help or oppose dissolving, and temperature tips the balance through the TΔS term. Enthalpy of solution (ΔH) Breaking lattice + solute–solvent attractions (hydration) compete. Can be endo- or exothermic depending on which dominates. Cold packs: endothermic dissolving. Entropy of solution (ΔS) Usually POSITIVE — ions spread out from an ordered lattice into solution. This positive ΔS often drives endothermic salts to dissolve anyway. TΔS grows with temperature. Dissolving is governed by ΔG = ΔH − TΔS — positive ΔS often drives even endothermic salts in. The Review Hub · AP Chemistry Unit 9 TOPIC 9.7 Coupled Reactions Pairing an unfavorable step with a favorable one An unfavorable reaction (ΔG > 0) can be driven by COUPLING it to a strongly favorable one (ΔG ≪ 0), as long as they share a common intermediate. The ΔG values ADD; the total must be negative. How coupling works Rxn 1: A → B, ΔG = +20 kJ Rxn 2: B → C, ΔG = −50 kJ Overall A → C: ΔG = −30 kJ → now FAVORABLE. B is the shared intermediate. Real-world coupling Living cells couple unfavorable reactions to ATP hydrolysis (ΔG ≪ 0) to drive biosynthesis. Metallurgy couples ore reduction to CO → CO₂ oxidation. Couple reactions so the ΔG values add to a negative total via a shared intermediate. The Review Hub · AP Chemistry Unit 9 TOPIC 9.8 Galvanic and Electrolytic Cells GALVANIC (VOLTAIC) CELL V e⁻ → salt bridge ANODE (−) oxidation CATHODE (+) reduction Galvanic: spontaneous ΔG < 0, E°_cell > 0. A favorable redox reaction generates electricity — this is a battery. Electrons flow anode → cathode. Electrolytic: driven ΔG > 0, E°_cell < 0. An external power source forces a nonspontaneous reaction. Both use "an ox, red cat": ANode = OXidation, REDuction at CAThode. Galvanic = spontaneous (E° > 0); electrolytic = driven (E° < 0) — oxidation always at the anode. The Review Hub · AP Chemistry Unit 9 TOPIC 9.9 Cell Potential and Free Energy ΔG° = −nFE°_cell The symbols n = moles of electrons transferred F = Faraday's constant, 96,485 C/mol E°_cell = standard cell potential (V) E°_cell = E°_cathode − E°_anode Use standard reduction potentials for both, then subtract (do not flip signs of E°). Sign relationship E°_cell > 0 ⇒ ΔG° < 0 ⇒ spontaneous (galvanic). E°_cell < 0 ⇒ ΔG° > 0 ⇒ nonspontaneous (electrolytic). A bigger positive E° means a stronger driving force (more negative ΔG°). ΔG° = −nFE° — positive E° ⇒ negative ΔG° ⇒ spontaneous cell. The Review Hub · AP Chemistry Unit 9 TOPIC 9.10 Cell Potential Under Nonstandard Conditions Concentration shifts the cell potential (Q vs. K) Away from standard conditions, E depends on the reaction quotient Q. As the cell runs, reactants deplete and products build up, so Q rises toward K and E falls — reaching E = 0 (dead battery) at equilibrium. Qualitative (Le Châtelier) view ↑ reactant conc. (or ↓ product): Q < K → E > E° (bigger push). ↑ product conc. (or ↓ reactant): Q > K → E < E° (smaller push). At equilibrium Q = K and E = 0. Nernst-style relationship E = E° − (RT/nF) ln Q. Larger Q lowers E; smaller Q raises it. A concentration cell (same electrodes, different concentrations) has E° = 0 but still produces voltage until Q = 1. Concentration changes E: Q < K raises it, Q > K lowers it, E = 0 at equilibrium. The Review Hub · AP Chemistry Unit 9 TOPIC 9.11 Electrolysis and Faraday's Law Charge → moles of electrons → moles of product In electrolysis, the amount of substance produced is proportional to the charge passed. Track it with a stoichiometry chain: current × time gives charge, then convert through the balanced half-reaction. q = I·t (C) → ÷ F = mol e⁻ → ÷ n = mol product → × M = grams Example: 2.0 A for 30 min q = 2.0 × 1800 = 3600 C → 3600 / 96,485 = 0.0373 mol e⁻. For Cu²⁺ + 2e⁻ → Cu: 0.0187 mol Cu = 1.19 g. Faraday's law: q = I·t → mol e⁻ (÷F) → mol product (÷n) → mass. The Review Hub · AP Chemistry Unit 9
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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 9's thermodynamics and electrochemistry content.