Review AP Chem Unit 9 to understand how entropy, Gibbs free energy, and electrochemistry connect energy changes to the direction and extent of chemical processes. This unit ties together thermodynamic favorability, equilibrium, and redox reactions in electrochemical cells.
Fiveable's AP Chemistry Unit 9 review covers 11 topics, with guided practice questions for the unit.
Use the topic guides, key terms, and practice questions available for every topic in this unit to build a complete review.
Unit 9 asks you to predict and explain energy changes in chemical systems using two connected frameworks: thermodynamics and electrochemistry. You will calculate entropy changes, evaluate thermodynamic favorability using Gibbs free energy, connect ΔG° to equilibrium constants, and then apply all of that reasoning to galvanic and electrolytic cells.
Unit 9 is about predicting whether chemical processes are thermodynamically favored (ΔG° < 0), how that connects to equilibrium constants and cell potentials, and how to use Faraday's law to relate charge, current, and mass in electrochemical cells.
Entropy (S) measures dispersal of matter and energy. The sign of ΔS° tells you whether disorder increases or decreases. Gibbs free energy combines enthalpy and entropy: ΔG° = ΔH° - TΔS°. When ΔG° < 0, the process is thermodynamically favored. Temperature determines which sign of ΔH° and ΔS° combinations are favored.
ΔG° and K are linked by ΔG° = -RT ln K. A negative ΔG° means K > 1 and products are favored at equilibrium. Thermodynamically unfavorable processes can be driven by coupling them to favorable reactions (shared intermediates) or by supplying external energy such as electricity or light.
Galvanic cells convert thermodynamically favored redox reactions into electrical energy; electrolytic cells use external power to drive unfavorable reactions. Standard cell potential E°cell = E°cathode - E°anode, and ΔG° = -nFE°. The Nernst equation describes how cell potential changes under nonstandard conditions, and Faraday's law connects charge, current, time, and mass deposited.
Every topic in Unit 9 returns to one question: will this process occur, and how much energy is involved? Entropy measures dispersal, Gibbs free energy combines enthalpy and entropy into a single favorability criterion, and electrochemistry translates that criterion into measurable voltage and calculable mass. The equation ΔG° = -nFE° = -RT ln K is the backbone of the entire unit, linking thermodynamics, equilibrium, and electrochemistry into one coherent framework.
Predict the sign and relative magnitude of ΔS for phase changes, volume changes, and reactions based on dispersal of matter and energy and changes in moles of gas.
Calculate ΔS°reaction using ΔS°reaction = ΣS°products - ΣS°reactants with tabulated standard molar entropy values. Remember units are J/(mol·K) and physical states matter.
Use ΔG° = ΔH° - TΔS° or ΔG°reaction = ΣΔG°f products - ΣΔG°f reactants to determine whether a process is thermodynamically favored. ΔG° < 0 means favored.
A thermodynamically favored reaction (ΔG° < 0) may not occur at a measurable rate if activation energy is high. Distinguish thermodynamic favorability from reaction rate in your explanations.
Connect ΔG° to K using ΔG° = -RT ln K. Negative ΔG° means K > 1 (products favored); positive ΔG° means K < 1 (reactants favored). Estimate K magnitude qualitatively from the size of ΔG° relative to RT.
Explain salt solubility using the three contributions to ΔG° of dissolution: lattice breaking (endothermic), solvent reorganization, and ion-dipole interactions. Predictions are difficult because these contributions partially cancel.
Drive unfavorable reactions by coupling them to favorable ones through a shared intermediate. Add ΔG° values to determine overall favorability. External energy sources (electricity, light) can also drive unfavorable processes.
Identify the anode (oxidation) and cathode (reduction) in both cell types. Galvanic cells produce energy (ΔG° < 0); electrolytic cells consume energy (ΔG° > 0). Track electron flow, ion flow through the salt bridge, and electrode mass changes.
Calculate E°cell = E°cathode - E°anode using standard reduction potentials. Connect to thermodynamics using ΔG° = -nFE°. Positive E°cell means thermodynamically favored; negative means unfavored.
Cell potential changes when concentrations deviate from standard conditions. The cell moves toward equilibrium as it operates; E approaches zero when Q = K. Reason qualitatively using Q relative to 1 and K.
Use I = q/t and the Faraday constant (96,485 C/mol e-) to convert between current, time, charge, moles of electrons, and mass deposited or removed at an electrode. Set up the dimensional analysis chain carefully.
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Review Gibbs Free Energy and Thermodynamic Favorability with attention to how the concept appears in AP-style source and evidence questions.
Review Cell Potential and Free Energy with attention to how the concept appears in AP-style source and evidence questions.
Review Galvanic (Voltaic) and Electrolytic Cells with attention to how the concept appears in AP-style source and evidence questions.
Review Absolute Entropy and Entropy Change with attention to how the concept appears in AP-style source and evidence questions.
Entropy (S) measures how dispersed matter and energy are in a system. You need to predict the sign and relative magnitude of ΔS for physical and chemical processes without calculating exact values. Focus on the direction of change: does the process increase or decrease disorder?
For the reaction 2H2(g) + O2(g) → 2H2O(g), predict the sign of ΔS° and explain your reasoning using moles of gas.
| Process | ΔS sign | Reason |
|---|---|---|
| Solid melting to liquid | Positive | Particles gain translational freedom |
| Liquid vaporizing to gas | Positive (large) | Large increase in volume and freedom |
| Gas compressed to smaller volume | Negative | Fewer accessible positions |
| 3 mol gas → 1 mol gas (reaction) | Negative | Fewer moles of gas means fewer microstates |
| Increasing temperature of a gas | Positive | Broader kinetic energy distribution |
Unlike enthalpy, entropy has a true absolute value anchored at zero for a perfect crystal at 0 K (Third Law). Standard molar entropy values S° are tabulated and used to calculate ΔS° for reactions using a products-minus-reactants summation, just like Hess's law for enthalpy.
Using tabulated S° values, calculate ΔS° for N2(g) + 3H2(g) → 2NH3(g) and explain whether the sign matches your prediction from moles of gas.
Gibbs free energy (ΔG°) combines enthalpy and entropy into a single criterion for thermodynamic favorability. A negative ΔG° means the process is thermodynamically favored under standard conditions. However, a favorable ΔG° does not guarantee the reaction occurs at a measurable rate. When a high activation energy prevents a favorable reaction from proceeding, the system is under kinetic control.
A reaction has ΔH° = +50 kJ/mol and ΔS° = +200 J/(mol·K). At what temperature does the reaction become thermodynamically favored? If the reaction still does not occur at that temperature, what explains this?
| ΔH° | ΔS° | Favored at... |
|---|---|---|
| Negative | Positive | All temperatures |
| Positive | Negative | No temperature (never favored) |
| Negative | Negative | Low temperatures only |
| Positive | Positive | High temperatures only |
ΔG° and K describe the same equilibrium position from two different angles. A large negative ΔG° corresponds to a large K (products strongly favored); a large positive ΔG° corresponds to a small K (reactants strongly favored). Dissolution applies this framework to salts, where ΔG° of dissolution reflects competing enthalpy and entropy contributions.
A reaction has ΔG° = -17.1 kJ/mol at 298 K. Estimate K qualitatively and then calculate it using ΔG° = -RT ln K.
A thermodynamically unfavorable reaction (ΔG° > 0) can be driven by coupling it to a favorable reaction (ΔG° < 0) that shares a common intermediate. The total ΔG° is the sum of the individual ΔG° values. If the sum is negative, the overall coupled process is thermodynamically favored.
Reaction A has ΔG° = +25 kJ/mol and reaction B has ΔG° = -40 kJ/mol. If they share a common intermediate, is the coupled process thermodynamically favored? What is the overall ΔG°?
Galvanic cells use thermodynamically favored redox reactions to generate voltage. Electrolytic cells use an external power source to drive unfavorable reactions. In both cell types, oxidation occurs at the anode and reduction occurs at the cathode. Standard cell potential E°cell connects directly to ΔG° through ΔG° = -nFE°.
For a Zn-Cu galvanic cell (E°(Zn2+/Zn) = -0.76 V, E°(Cu2+/Cu) = +0.34 V), identify the anode and cathode, calculate E°cell, and determine ΔG° for the transfer of 2 moles of electrons.
| Feature | Galvanic Cell | Electrolytic Cell |
|---|---|---|
| ΔG° | Negative | Positive |
| E°cell | Positive | Negative |
| Energy flow | Produces electrical energy | Consumes electrical energy |
| Anode charge | Negative | Positive |
| Example | Zn-Cu Daniell cell | Electroplating, charging a battery |
Under nonstandard conditions, cell potential shifts from E°cell depending on the reaction quotient Q. The cell is always moving toward equilibrium, so E approaches zero as Q approaches K. Faraday's law connects the charge passed through an electrolytic cell to the mass of material deposited or removed at an electrode.
A current of 2.00 A is passed through a solution of CuSO4 for 965 seconds. How many grams of copper are deposited at the cathode? (Cu2+ + 2e- → Cu, M = 63.55 g/mol)
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| Term | Definition |
|---|---|
| ΔG° = ΔH° − TΔS° | The fundamental equation relating Gibbs free energy change to enthalpy change, entropy change, and absolute temperature. When ΔG° < 0, the process is thermodynamically favored. |
| ΔG° = -RT ln K | Relates standard free energy change to the equilibrium constant. Negative ΔG° gives K > 1 (products favored); positive ΔG° gives K < 1 (reactants favored). |
| ΔG° = -nFE° | Connects standard Gibbs free energy change to standard cell potential, where n is moles of electrons transferred and F is the Faraday constant (96,485 C/mol). |
| ΔS°reaction = ΣS°products - ΣS°reactants | Calculates standard entropy change of a reaction from tabulated standard molar entropy values. Units are J/(mol·K). |
| microstates | The number of distinct molecular arrangements available to a system. Entropy increases as the number of accessible microstates increases. |
| Standard Cell Potential | E°cell = E°cathode - E°anode, calculated from standard reduction potentials. A positive E°cell indicates a thermodynamically favored redox reaction. |
| Galvanic Cell | An electrochemical cell that converts a thermodynamically favored redox reaction (ΔG° < 0, E°cell > 0) into electrical energy. Also called a voltaic cell. |
| Electrolytic Cell | An electrochemical cell that uses an external power source to drive a thermodynamically unfavorable redox reaction (ΔG° > 0, E°cell < 0). |
| Anode | The electrode where oxidation occurs in any electrochemical cell. Electrons leave the cell through the anode into the external circuit. |
| Cathode | The electrode where reduction occurs in any electrochemical cell. Electrons arrive at the cathode from the external circuit. |
| Faraday constant | 96,485 coulombs per mole of electrons. Used to convert between charge passed and moles of electrons in Faraday's law calculations. |
| I = q/t | Current (amperes) equals charge (coulombs) divided by time (seconds). The starting equation for all Faraday's law stoichiometry problems. |
| Nonstandard Conditions | Any conditions where concentrations, pressures, or temperature differ from standard state (1 M, 1 atm, 298 K). Cell potential shifts from E°cell based on the reaction quotient Q. |
| Activation Energy | The energy barrier a reaction must overcome to proceed. A high activation energy can prevent a thermodynamically favored reaction from occurring at a measurable rate, placing it under kinetic control. |
| Salt Bridge | Connects the two half-cells of a galvanic cell, allowing ion flow to maintain electrical neutrality without mixing the electrode solutions. |
A negative ΔG° means a process is thermodynamically favored, not that it is fast. Diamond converting to graphite has ΔG° < 0 but does not occur at a measurable rate at room temperature because of a very high activation energy. Always separate these two ideas in your explanations.
Standard molar entropy S° is in J/(mol·K), but ΔH° and ΔG°f are typically in kJ/mol. When using ΔG° = ΔH° - TΔS°, convert ΔS° to kJ/(mol·K) by dividing by 1000 before substituting.
Electrochemical cells are not at equilibrium while operating, so Le Chatelier's principle does not apply. Use Q vs. K reasoning to explain how concentration changes affect cell potential.
In galvanic cells the anode is negative; in electrolytic cells the anode is positive (connected to the positive terminal of the power supply). The rule that oxidation occurs at the anode and reduction at the cathode never changes, but the charge on the electrode does.
The number of electrons transferred (n) comes from the balanced half-reaction, not from the overall reaction. For Cu2+ + 2e- → Cu, n = 2 per mole of copper. Using n = 1 here would double the calculated mass.
AP Chemistry questions frequently ask you to explain whether a process is thermodynamically favored and to support your answer using ΔG°, ΔH°, ΔS°, K, or E°cell. Practice connecting these quantities explicitly: state the sign of ΔG°, explain which enthalpy and entropy contributions produce that sign, and confirm with K > 1 or E°cell > 0 as appropriate. A complete justification uses at least two connected pieces of evidence.
A common task in AP Chemistry is to explain why a thermodynamically favored process does not occur at a measurable rate. Your response must name activation energy as the kinetic barrier and explicitly state that ΔG° < 0 does not determine rate. Mixing up these two concepts in a written explanation is a frequent source of lost points.
Electrochemistry problems often chain several calculations: identify half-reactions and assign anode/cathode, calculate E°cell, convert to ΔG° using ΔG° = -nFE°, and then apply Faraday's law to find mass deposited or time required. On the AP exam, show each conversion step clearly and include units throughout. Errors in identifying n (moles of electrons from the balanced half-reaction) are the most common source of incorrect answers in these problems.
For any process, identify whether matter or energy becomes more or less dispersed. Apply the moles-of-gas rule for reactions and the phase-change rule for physical processes.
Use ΔS°reaction = ΣS°products - ΣS°reactants and ΔG°reaction = ΣΔG°f products - ΣΔG°f reactants. Watch units: S° is in J/(mol·K) and ΔG°f is in kJ/mol.
Use the four-case sign table (ΔH° and ΔS° combinations) to determine at which temperatures a process is thermodynamically favored. Identify the crossover temperature when ΔG° = 0.
Given ΔG°, determine whether K > 1 or K < 1 and estimate its magnitude. Given K, determine the sign and approximate magnitude of ΔG°.
For any cell, identify the anode and cathode, the direction of electron flow, the role of the salt bridge, and whether the cell is galvanic or electrolytic. Calculate E°cell and ΔG° using ΔG° = -nFE°.
Explain qualitatively how Q relative to 1 and K affects cell potential. Know that E = 0 at equilibrium and that Le Chatelier's principle does not apply to electrochemical systems.
Chain the calculation: current and time to charge (q = I x t), charge to moles of electrons (divide by 96,485 C/mol), moles of electrons to moles of substance via half-reaction stoichiometry, then to mass.
Read the topic guides for 9.1 and 9.2. Practice predicting ΔS° sign for phase changes, gas expansions, and reactions by counting moles of gas. Then practice the ΔS°reaction = ΣS°products - ΣS°reactants calculation with tabulated values, paying attention to physical states and units.
Use the four-case ΔH°/ΔS° sign table to predict thermodynamic favorability at different temperatures. Practice calculating ΔG° both from ΔG° = ΔH° - TΔS° and from formation values. Then practice writing explanations that separate thermodynamic favorability from kinetic control using activation energy.
Practice converting between ΔG° and K using ΔG° = -RT ln K. Focus on qualitative estimation: large negative ΔG° means very large K. Apply this framework to dissolution problems, identifying the enthalpy and entropy contributions to ΔG° of dissolution and explaining why predictions can be difficult.
Practice adding ΔG° values for coupled reactions and identifying common intermediates. Then work through galvanic and electrolytic cell diagrams: label anode, cathode, electron flow, and salt bridge. Calculate E°cell from standard reduction potentials and convert to ΔG° using ΔG° = -nFE°.
Practice qualitative Nernst reasoning: determine whether a concentration change increases or decreases |E| relative to E°cell. Then work Faraday's law problems from start to finish using the full dimensional analysis chain from current and time to mass deposited. Use available practice questions to check your setup before solving.
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get printablesAP Chem Unit 9 covers 11 topics across thermodynamics and electrochemistry: entropy (9.1, 9.2), Gibbs free energy and thermodynamic favorability (9.3), thermodynamic vs. kinetic control (9.4), free energy and equilibrium (9.5), free energy of dissolution (9.6), coupled reactions (9.7), galvanic and electrolytic cells (9.8), cell potential and free energy (9.9), cell potential under nonstandard conditions (9.10), and electrolysis and Faraday's Law (9.11). The unit connects energy changes at the molecular level to macroscopic outcomes, so you'll see how entropy, Gibbs free energy, and electrochemical cells all tie together. Check out AP Chem Unit 9 for topic-by-topic breakdowns.
AP Chem Unit 9 makes up 7-9% of the AP exam. That weight covers thermodynamics and electrochemistry, including entropy, Gibbs free energy, equilibrium relationships, galvanic and electrolytic cells, and Faraday's Law. It's a focused unit, but the concepts show up in calculation-heavy multiple-choice and free-response questions. Because the percentage is on the smaller side, students sometimes underestimate this unit. The math-intensive topics like cell potential and Gibbs free energy calculations tend to appear on the FRQ section, so solid practice here pays off.
The AP Chem Unit 9 progress check includes both MCQ and FRQ parts drawn from all 11 topics in the unit. The MCQ section tests conceptual understanding of entropy, Gibbs free energy, thermodynamic favorability, and cell potential. The FRQ part typically asks you to calculate delta-G, interpret equilibrium relationships using free energy, or analyze electrochemical cells including electrolysis. For the progress check FRQ, expect to show your work on Gibbs free energy calculations, connect free energy to equilibrium constants, and explain how electrolysis and Faraday's Law apply to a given scenario. Practicing those question types on AP Chem Unit 9 before you submit the progress check is a smart move.
AP Chem Unit 9 FRQs most often focus on Gibbs free energy calculations, the relationship between free energy and equilibrium, cell potential under nonstandard conditions, and electrolysis with Faraday's Law. To practice, work through problems that ask you to calculate delta-G from delta-H and delta-S, connect delta-G to the equilibrium constant K, and determine the amount of substance produced during electrolysis. A few tips that help:
- Write out every step of your calculation, since partial credit is awarded for correct setup even if the final answer is wrong.
- Practice interpreting the sign of delta-G to predict thermodynamic favorability.
- For electrochemistry FRQs, make sure you can draw and label a galvanic cell and explain the direction of electron flow. You can find FRQ-style practice matched to these topics at AP Chem Unit 9.
For AP Chem Unit 9 multiple-choice and practice test questions, AP Chem Unit 9 is the best starting point, with MCQ and FRQ practice organized by topic across all 11 topics in the unit. Look for questions covering entropy, Gibbs free energy, electrolysis, and cell potential, since those are the highest-yield areas for both MCQ and the full practice test. When you work through MCQs, focus on questions that ask you to predict thermodynamic favorability, interpret free energy and equilibrium relationships, and calculate cell potential. Mixing conceptual MCQs with calculation-based practice gives you the best coverage of what shows up on the real exam.
Start AP Chem Unit 9 by building a strong foundation in entropy before moving to Gibbs free energy, since delta-G ties together delta-H, delta-S, and temperature in one equation you'll use constantly. Once that relationship clicks, connecting free energy to equilibrium constants and cell potential becomes much more straightforward. A concrete study plan that works:
1. Learn the entropy rules (9.1-9.2): predict whether delta-S is positive or negative from the reaction.
2. Practice Gibbs free energy calculations (9.3) until the sign conventions feel automatic.
3. Work through the free energy and equilibrium connection (9.5) using real K values.
4. Study galvanic vs. electrolytic cells (9.8) side by side so you don't mix them up.
5. Finish with electrolysis and Faraday's Law (9.11), which is very calculation-driven. Review topic by topic at AP Chem Unit 9, then test yourself with progress check-style questions to find gaps before the exam.