What it covers: Energy and change — endothermic/exothermic processes, calorimetry, phase change energy, bond enthalpies, enthalpy of formation, and Hess's law.
Exam weight: About 7–9% of the AP Chemistry exam.
The big question: How can the energy change of a chemical or physical process be measured, predicted, or calculated using multiple independent methods?
Big Ideas covered: Transformations (TRA), Scale, Proportion & Quantity (SPQ).
Key topics at a glance
Endothermic & Exothermic
Endothermic = absorbs energy (+ΔH). Exothermic = releases energy (−ΔH).
Heat Transfer
Heat flows hot → cold until thermal equilibrium. Heat (q) ≠ temperature.
Calorimetry
q = mcΔT. Use specific heat capacity to relate heat, mass, and temperature change.
Phase Changes
Energy goes into overcoming IMFs, not raising temperature — that's why temperature plateaus during melting/boiling.
Enthalpy (H)
A state function — depends only on initial/final states, not the path. ΔH at constant pressure = heat released/absorbed.
Bond Enthalpy Method
ΔHrxn ≈ (energy to break reactant bonds) − (energy to form product bonds).
Enthalpy of Formation Method
ΔHrxn = Σ ΔHf°(products) − Σ ΔHf°(reactants). Elements in standard state have ΔHf° = 0.
Hess's Law
Add/reverse/scale known reactions to build the target reaction; ΔH values add the same way the equations do.
The key terms you must know
Endothermic / exothermic — process absorbs energy (+ΔH) vs. releases energy (−ΔH).
System / surroundings — what's being studied vs. everything else exchanging energy with it.
q = mcΔT — the core calorimetry equation.
Heat of fusion / vaporization — energy for phase changes, with no temperature change.
Enthalpy (H) / state function — depends only on initial and final states, not the path.
Bond enthalpy — energy to break a specific covalent bond; used to estimate ΔHrxn.
Standard enthalpy of formation (ΔHf°) — enthalpy to form 1 mol of compound from elements in standard state (zero for elements).
Hess's law — combine known reactions to find ΔH for an unmeasured reaction.
Key themes to remember
Enthalpy is a state function — that's the whole unit in one sentence. Every calculation method (bond enthalpy, formation, Hess's law) exists because the path doesn't matter, only start and end points.
Heat and temperature are not the same thing. Heat is energy transferred; temperature is a measure of average kinetic energy.
Phase changes "hide" energy. Temperature plateaus during melting/boiling because all the added energy is overcoming intermolecular forces, not increasing kinetic energy.
Sign conventions matter enormously. Get comfortable flipping the sign of ΔH when you reverse a reaction in a Hess's law problem.
This unit sets up Unit 9. Enthalpy as a state function returns when Gibbs free energy (ΔG = ΔH − TΔS) combines it with entropy.
Common exam traps
Reversing a reaction flips the sign of ΔH. Forgetting this is the single most common Hess's law error.
Scaling a reaction by a coefficient scales ΔH by the same factor. Doubling a reaction's coefficients doubles its ΔH.
Bond enthalpy calculations always subtract products from reactants in the wrong order if you're not careful. ΔHrxn = Σ(bonds broken in reactants) − Σ(bonds formed in products) — breaking bonds costs energy (positive), forming bonds releases energy (already accounted for by subtracting).
An element in its standard state has ΔHf° = 0 — but only in its standard state. O₂(g) has ΔHf° = 0, but O₃(g) does not.
Don't confuse exothermic with "no activation energy needed." Exothermic reactions can still have a large activation energy barrier (see Unit 5) even though the overall reaction releases energy.