AP Chemistry Unit 6: Thermodynamics
Study enthalpy, entropy, Gibbs free energy, Hess's law, calorimetry with exam-format practice and rubric-based scoring.
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Inside This Unit: The Full Breakdown
Thermodynamics covers the energy changes in chemical reactions and physical processes. This unit focuses on enthalpy, calorimetry, Hess's law, and bond energies — the tools for calculating how much heat a reaction absorbs or releases.
Why it matters
Thermodynamics calculations are a guaranteed part of the AP Chemistry exam. You must perform calorimetry calculations, apply Hess's law, and use standard enthalpies of formation. These quantitative skills are tested in both multiple choice and free response.
Key concepts
- Enthalpy (ΔH) is the heat change at constant pressure. Exothermic reactions have ΔH < 0 (release heat); endothermic reactions have ΔH > 0 (absorb heat).
- Calorimetry measures heat transfer using q = mcΔT for specific heat or q = nΔH for molar enthalpy changes.
- Hess's law: the total enthalpy change for a reaction is the same regardless of the pathway — you can add ΔH values for individual steps.
- Standard enthalpy of formation (ΔH°f) allows calculation of reaction enthalpy: ΔH°rxn = Σ ΔH°f(products) − Σ ΔH°f(reactants).
Enthalpy and Calorimetry
Enthalpy (H) is a measure of heat content at constant pressure. The enthalpy change (ΔH) for a reaction equals the heat absorbed or released. Exothermic reactions release heat (ΔH < 0) and feel warm; endothermic reactions absorb heat (ΔH > 0) and feel cold. Calorimetry measures these changes experimentally: in a coffee-cup calorimeter, q = mcΔT, where m is the mass of solution, c is the specific heat capacity (4.18 J/g·°C for water), and ΔT is the temperature change. The heat absorbed by the solution equals the heat released by the reaction (and vice versa), with opposite sign. State functions like enthalpy depend only on the initial and final states, not the path — this is why Hess's law works.
Hess's Law and Enthalpies of Formation
Hess's law states that enthalpy is a state function, so the total ΔH for a reaction is the same whether it occurs in one step or multiple steps. If you can write a target reaction as the sum of known reactions, the ΔH values add up accordingly. When reversing a reaction, change the sign of ΔH. When multiplying a reaction by a coefficient, multiply ΔH by the same factor. Standard enthalpies of formation (ΔH°f) provide a systematic alternative: ΔH°rxn = Σ ΔH°f(products) − Σ ΔH°f(reactants). By definition, ΔH°f for any element in its standard state is zero. This formula is one of the most useful on the entire AP exam and works for any reaction when formation data is available.
Bond Energies
Bond energy is the energy required to break one mole of a bond in the gas phase. Bond breaking is endothermic (requires energy input), and bond forming is exothermic (releases energy). The enthalpy of a reaction can be estimated from bond energies: ΔH ≈ Σ(bonds broken) − Σ(bonds formed). This method gives approximate values because bond energies are averages that vary slightly depending on molecular context. In exothermic reactions, the bonds formed in the products are stronger (more energy released) than the bonds broken in the reactants. Bond energy calculations are particularly useful when standard formation data is unavailable. Note that this method works best for gas-phase reactions and gives less accurate results for reactions involving liquids or solids.
AP exam tip
On the AP exam, be very careful with signs. Heat lost by the system is gained by the surroundings. If ΔT is positive in a calorimetry experiment, the reaction is exothermic (ΔH is negative) because the reaction released heat into the water.
Connections to other units
- Unit 3 (Intermolecular Forces): Phase changes involve enthalpy changes associated with overcoming IMFs.
- Unit 5 (Kinetics): Activation energy is related to bond breaking; thermodynamics determines overall favorability.
- Unit 9 (Applications of Thermodynamics): Free energy combines enthalpy and entropy to predict spontaneity.
