What it covers: Entropy, Gibbs free energy, thermodynamic favorability, and how free energy connects to equilibrium and electrochemistry (galvanic and electrolytic cells).
Exam weight: About 7–9% of the AP Chemistry exam — one of the smaller units, but conceptually dense.
The big question: How does Gibbs free energy unify thermodynamics, equilibrium, and electrochemistry into one consistent framework?
Big Ideas covered: Energy (ENE), Transformations (TRA).
Key topics at a glance
Entropy (S)
A measure of disorder/microstates. The second law: total entropy of the universe increases for any spontaneous process.
Entropy Change (ΔS)
Positive ΔS when moles of gas increase, a substance melts/boils, or disorder increases overall.
Gibbs Free Energy
ΔG = ΔH − TΔS. Negative ΔG means a process is thermodynamically favorable at that temperature.
Thermodynamic vs. Kinetic Control
Favorable (ΔG < 0) ≠ fast. Kinetics (activation energy) determines rate, separate from thermodynamics.
Free Energy & Equilibrium
ΔG° = −RT ln K. A more negative ΔG° corresponds to a larger equilibrium constant K.
Coupled Reactions
Pairing an unfavorable reaction with a favorable one so their combined ΔG is negative, driving the unfavorable one forward.
Standard reduction potential (E°) — a half-reaction's tendency to be reduced, relative to the standard hydrogen electrode.
ΔG = −nFE — links free energy to cell potential.
Nernst equation — calculates cell potential under nonstandard concentrations.
Key themes to remember
Temperature can flip the favorability of a reaction. A process with unfavorable ΔH but favorable ΔS (or vice versa) can switch from favorable to unfavorable depending on temperature.
Thermodynamic favorability and reaction rate are completely independent ideas. A favorable reaction (ΔG < 0) can still be extremely slow if its activation energy is high.
ΔG, K, and E°cell are three views of the same underlying chemistry. A very negative ΔG° means a large K AND a large positive E°cell — they all move together.
Galvanic cells generate electricity; electrolytic cells consume it. The labels for anode (oxidation) and cathode (reduction) stay the same in both — only the direction of spontaneous electron flow differs.
This unit is the capstone that ties Units 6, 7/8, and electrochemistry together. Every equation here connects back to enthalpy, entropy, or equilibrium concepts you've already learned.
Common exam traps
Don't confuse "thermodynamically favorable" with "fast." ΔG < 0 says nothing about reaction rate — that's a kinetics question.
Watch your signs carefully in ΔG = ΔH − TΔS. A common error is forgetting that increasing T makes the −TΔS term more negative when ΔS is positive (more favorable), and more positive when ΔS is negative (less favorable).
Don't mix up the anode and cathode between galvanic and electrolytic cells. Oxidation always happens at the anode and reduction always happens at the cathode — but which electrode is positive/negative flips between the two cell types.
Remember units consistency between ΔH (kJ) and ΔS (J/K) in ΔG = ΔH − TΔS. A common arithmetic mistake is forgetting to convert one of these before combining them.
Standard reduction potentials (E°) are NOT additive the way ΔG values are when reversing a half-reaction. E° for an oxidation is just the negative of the reduction E°, but you don't multiply E° by stoichiometric coefficients the way you would ΔG.