⚗️
AP Chemistry · Unit 6 of 9
Thermodynamics
7–9% of the exam6 lessons · 82 min18 terms
What this unit covers
The topics below follow the published Chemistry course framework for Unit 6. This unit is worth 7–9% of the exam, so budget your time against that rather than against how long the unit takes to teach.
EnthalpyCalorimetryHess’s lawEntropy & Gibbs free energy
Lessons in this unit
- Energy, Heat & Enthalpy12 min · 3 objectivesDistinguish endothermic from exothermic changes and read their energy diagrams · Separate the ideas of heat, temperature, and enthalpy · Interpret the sign of ΔH and scale it with the amount of reactant
- Calorimetry13 min · 3 objectivesUse q = mcΔT to relate heat, mass, specific heat, and temperature change · Apply conservation of energy between system and surroundings in a calorimeter · Reason about how specific heat controls a substance's temperature response
- Hess's Law & Enthalpy of Formation14 min · 3 objectivesApply Hess's law by adding, reversing, and scaling thermochemical equations · Compute ΔH°rxn from standard enthalpies of formation · Estimate ΔH from bond enthalpies (bonds broken minus bonds formed)
- Entropy & Gibbs Free Energy14 min · 3 objectivesPredict the sign of ΔS from changes in phase, moles of gas, and disorder · Combine ΔH and ΔS into ΔG = ΔH − TΔS to judge spontaneity · Analyze how the signs of ΔH and ΔS set the temperature range of spontaneity
- Calorimetry & Hess's Law — Advanced Problem-Solving15 min · 3 objectivesSolve two-substance calorimetry problems with conservation of energy to find a final temperature · Analyze bomb-calorimeter data using a calorimeter heat capacity to report ΔH per mole · Assemble reactions with Hess's law and cross-check ΔH°rxn from formation and bond enthalpies
- Entropy & Gibbs Free Energy — Spontaneity & Temperature14 min · 3 objectivesPredict the sign of ΔS°rxn from phase changes, gas-mole counts, and standard entropies · Calculate ΔG from ΔH and ΔS and classify the four sign combinations by temperature dependence · Solve for the crossover temperature where ΔG = 0 and spontaneity flips
Formulas in Unit 6
Sign convention for enthalpy
ΔH = H(products) − H(reactants); exothermic ΔH < 0, endothermic ΔH > 0
Negative means energy leaves the system (released); positive means energy enters the system (absorbed).
Heat and temperature change
q = m·c·ΔT
q = heat (J), m = mass (g), c = specific heat capacity (J·g⁻¹·°C⁻¹), ΔT = T(final) − T(initial). For water, c = 4.18 J·g⁻¹·°C⁻¹.
Enthalpy from formation enthalpies
ΔH°rxn = Σ ΔH°f(products) − Σ ΔH°f(reactants)
Each ΔH°f is multiplied by its coefficient. The ΔH°f of any element in its standard state is exactly 0.
Gibbs free energy
ΔG = ΔH − T·ΔS
T is absolute temperature in kelvin. Watch units: ΔH is usually kJ while ΔS is usually J·K⁻¹ — convert ΔS to kJ·K⁻¹ (divide by 1000) before subtracting.
Two cross-checks for ΔH°rxn
ΔH°rxn = Σ ΔH°f(products) − Σ ΔH°f(reactants); ΔH ≈ Σ(bonds broken) − Σ(bonds formed)
Each ΔH°f is weighted by its coefficient; ΔH°f of an element in its standard state is 0. The bond-enthalpy estimate is gas-phase only and approximate (tabulated bonds are averages).
Entropy change of a reaction
ΔS°rxn = Σ S°(products) − Σ S°(reactants)
S° values are in J·mol⁻¹·K⁻¹ and are always positive (even for elements). More moles of gas on the product side → ΔS > 0.
Crossover (equilibrium) temperature
ΔG = 0 ⟹ T = ΔH / ΔS
Use ΔH in joules to match ΔS in J·K⁻¹ (or ΔH in kJ with ΔS in kJ·K⁻¹). For (+, +) the reaction turns spontaneous above T; for (−, −) it turns nonspontaneous above T.
Every term in Unit 6
All 18 terms we publish for Thermodynamics, with definitions. Reading them through is the fastest way to find the ones you cannot define — then drill those in cram mode until you can produce them without the prompt.
- Endothermic vs exothermic
- Endothermic absorbs heat, ΔH positive, surroundings cool. Exothermic releases heat, ΔH negative, surroundings warm.
- Heat capacity and q = mcΔT
- Energy transferred equals mass times specific heat times temperature change. Water's high specific heat is why it is the usual calorimetry medium.
- Calorimetry
- Heat lost by the system equals heat gained by the surroundings, assuming no loss to the environment — the assumption most error analysis targets.
- Enthalpy of reaction from bond energies
- ΔH = bonds broken minus bonds formed. Breaking absorbs energy and forming releases it.
- Hess's Law
- ΔH is a state function, so the enthalpy of an overall reaction is the sum of its steps. Reversing a step flips the sign; scaling it scales ΔH.
- Standard enthalpy of formation
- ΔH°f is the enthalpy change forming one mole of a compound from its elements in standard states. Elements in standard state have ΔH°f = 0.
- ΔH°rxn from formation enthalpies
- Sum of products' ΔH°f minus sum of reactants' ΔH°f, each multiplied by its coefficient.
- Entropy
- A measure of energy dispersal. Entropy rises going solid to liquid to gas, on dissolving, on heating, and when moles of gas increase.
- Predicting the sign of ΔS
- Count moles of gas on each side. More gas moles among the products means positive ΔS.
- Gibbs free energy
- ΔG = ΔH − TΔS. Negative ΔG means thermodynamically favorable at that temperature.
- Temperature dependence of ΔG
- ΔH negative with ΔS positive is favorable at all temperatures; the reverse is never favorable. Mixed signs make favourability temperature-dependent.
- ΔG and the equilibrium constant
- ΔG° = −RT ln K. A negative ΔG° means K > 1 and products are favored at equilibrium.
- Thermodynamic vs kinetic favourability
- A reaction can have a very negative ΔG and still not proceed measurably if its activation energy is high — diamond to graphite is the standard example.
- System vs surroundings
- The system is what you are studying; everything else is the surroundings. A negative ΔH means energy left the system and entered the surroundings.
- Coffee-cup calorimetry assumptions
- That no heat escapes and that the solution behaves like water. Both are approximations, and the usual source of error in the result.
- State function
- A quantity depending only on the current state, not the path taken. Enthalpy, entropy and free energy are state functions; work and heat are not.
- Why entropy of the universe matters
- A process is spontaneous when the total entropy of system plus surroundings increases. ΔG < 0 is that same criterion expressed using system properties only.
- Heating and cooling curves
- Sloped segments are temperature change (q = mcΔT); flat segments are phase change (q = mΔH) where energy goes to breaking intermolecular forces, not raising temperature.
What examiners penalize here
- A quick reflex for the exam: "exo = exit," heat exits the system, ΔH < 0, surroundings warm. "Endo = into," heat goes into the system, ΔH > 0, surroundings cool. The temperature change you feel in the beaker is the *opposite* sign of the system's ΔH.
- Free-response calorimetry problems almost always want the heat *of the reaction*, not just of the water. Compute q(water) with m·c·ΔT, then flip the sign to get q(reaction), and finally divide by moles to report ΔH in kJ·mol⁻¹.
- Three routes, one answer: Hess's law with given steps, the ΔH°f formula with a table, and bond enthalpies for gas-phase estimates. If a problem hands you ΔH°f values, reach for Σproducts − Σreactants first — it is the fastest and the most exam-common.
- Memorize the four-case table: (−, +) spontaneous at all T; (+, −) never; (−, −) spontaneous at low T; (+, +) spontaneous at high T. When ΔH and ΔS pull the same direction, temperature is irrelevant; when they conflict, T > ΔH/ΔS marks the switch point.
- On the free-response section, calorimetry usually feeds Hess or ΔH°f: measure q(reaction) with the calorimeter, divide by moles to get a molar ΔH, then use it as one step or cross-check it against Σ ΔH°f(products) − Σ ΔH°f(reactants). Always report the final sign — exothermic is negative.
- A reliable free-response routine: compute ΔH°rxn and ΔS°rxn from tables, decide the sign case, and if ΔH and ΔS conflict solve T = ΔH/ΔS for the crossover (keep units consistent — ΔH in J with ΔS in J·K⁻¹). Then answer the temperature question by naming which side of the crossover is spontaneous.
Practice Chemistry
Our practice bank is drawn from across the whole course rather than filtered to one unit, which is closer to how the exam asks anyway — it will not tell you which unit a question is testing.
Questions about this unit
How much of the AP Chemistry exam is Unit 6?
Unit 6, Thermodynamics, is worth 7–9% of the Chemistry multiple-choice section according to the published course framework. Across all 9 units that makes it a middling share, roughly what an even split across units would give.
What topics are covered in Chemistry Unit 6?
Thermodynamics covers Enthalpy, Calorimetry, Hess’s law and Entropy & Gibbs free energy. We publish 18 terms with definitions for this unit, all of them on this page.
How should I study Chemistry Unit 6?
Read the 6 lessons below first — about 80 minutes — then drill the 18 terms in cram mode until you can produce each definition from memory rather than just recognize it. Recognition is what makes a unit feel finished when it is not. Finish with practice questions and read the explanation for every one you get right by elimination as well as the ones you miss.
All 9 units of AP Chemistry
Unit names, topics and exam weights follow the published College Board course framework for AP Chemistry. AP® is a trademark registered by the College Board, which does not endorse this site.