AP Chemistry · Handsworth Secondary 2026–27
Unit 6 · Study Guide
Dr. Ras Mulinta
Thermochemistry
Exam-focused review
A one-page map of what the Unit 6 test (and the AP exam) expects. This is a checklist, not a re-teach, if a line doesn't click, go back to that section of the notes package. A periodic table and a ΔH°f / bond-energy table are provided on the test.
Must be able to do
- 6.1 Classify a process as endo- or exothermic from the temperature change of the surroundings, and give the sign of ΔH (exo < 0, endo > 0).
- 6.2 Draw and read an energy diagram: products lower = exothermic, products higher = endothermic; ΔH = E(products) − E(reactants).
- 6.3 Explain heat transfer at the particle level, warmer = higher average KE; collisions move energy until thermal equilibrium (equal temperatures).
- 6.4 Apply q = mcΔT and the calorimetry balance q_lost = −q_gained; solve for q, c, m, or final temperature.
- 6.5 Find phase-change energy with q = n·ΔH_phase; know melting/boiling absorb, freezing/condensing release the same magnitude.
- 6.6 Relate moles to reaction enthalpy: q = n·ΔH_rxn, scaling to the balanced equation.
- 6.7 Estimate ΔH from bonds: ΔH ≈ Σ(broken) − Σ(formed)count the bonds carefully.
- 6.8 Use formation data: ΔH°rxn = ΣΔH°f(products) − ΣΔH°f(reactants); elements in standard state = 0.
- 6.9 Combine reactions with Hess's law, reverse flips the sign, multiplying scales ΔH, adding sums ΔH.
The big idea that ties it together
Energy is conserved. Every tool in this unit, q = mcΔT, q = n·ΔH, bond enthalpies, formation tables, Hess's law, is the same first-law accounting: whatever energy the system loses, the surroundings gain, and vice versa. Track the sign (does the system release or absorb?), pick the matching equation, and the path you take doesn't change the total.
Don't waste time on (excluded by the CED)
The technical distinction between enthalpy and internal energy (treat ΔH as the heat of reaction at constant pressure) · the formal concept of state functions (just apply Hess's three rules) · entropy and Gibbs free energy (those belong to Unit 9, not here). Know the patterns, skip the edge cases.
Quick self-check (answer in your head, then verify)
- How much heat warms 100.0 g of water from 20.0 °C to 45.0 °C? (c = 4.18 J·g⁻¹·°C⁻¹)
- A salt dissolves and the beaker turns cold. Endo- or exothermic? Sign of ΔH?
- Energy to melt 2.00 mol of ice? (ΔH_fus = +6.01 kJ/mol)
- Estimate ΔH for H₂ + Cl₂ → 2 HCl. (H–H 436, Cl–Cl 242, H–Cl 431 kJ/mol)
- ΔH°rxn for 2 H₂O₂(l) → 2 H₂O(l) + O₂(g)? (ΔH°f: H₂O₂(l) −187.8, H₂O(l) −285.8, O₂ 0)
- Using N₂ + O₂ → 2 NO (ΔH = +180 kJ) and 2 NO + O₂ → 2 NO₂ (ΔH = −112 kJ), find ΔH for N₂ + 2 O₂ → 2 NO₂.
- On an energy diagram, reactants are at 120 kJ and products at 200 kJ. Find ΔH and classify it.
Check yourself: 1) q = (100.0)(4.18)(25.0) = 10450 J ≈ 1.05×10⁴ J (10.5 kJ) 2) endothermic, ΔH > 0 (system pulled energy from the surroundings) 3) q = (2.00)(6.01) = +12.0 kJ 4) (436+242) − 2(431) = 678 − 862 = −184 kJ (exothermic) 5) [2(−285.8)+0] − [2(−187.8)] = −571.6 + 375.6 = −196.0 kJ 6) (+180) + (−112) = +68 kJ 7) ΔH = 200 − 120 = +80 kJ, endothermic (products higher).