Unit 6: Thermodynamics
AP Chemistry: 81 practice questions with detailed explanations.
Unit Study Guide
Executive Summary
Unit 6 tracks energy flow. Enthalpy changes tell you whether a process releases or absorbs heat; Hess's law and calorimetry quantify it.
Endothermic vs. exothermic
Exothermic processes release energy (ΔH < 0, surroundings warm); endothermic processes absorb it (ΔH > 0). Bond breaking is endothermic; bond forming is exothermic. Energy diagrams plot reactants and products with Ea peaks.
Heat transfer and calorimetry
Heat flows from hot to cold until thermal equilibrium. q = mcΔT measures heat for a temperature change (c = specific heat, water 4.18 J/g·°C). Calorimetry: heat lost by reaction = heat gained by surroundings (coffee-cup calorimetry assumes constant pressure, q = ΔH).
Phase changes
During melting/vaporization, temperature is constant; energy breaks IMFs. q = nΔHfus or q = nΔHᵥap. Heating curves show plateaus at phase transitions; slopes involve c (specific heat) and q = mcΔT. ΔHᵥap of water ≈ 40.7 kJ/mol; ΔHfus ≈ 6.01 kJ/mol.
Enthalpy of reaction
ΔH°_rxn = Σ(products' ΔH°_f × coefficients) − Σ(reactants' ΔH°_f × coefficients). ΔH°_f is zero for elements in their standard states. Hess's law: ΔH for the sum of steps equals the sum of ΔH of the steps — reverse a step, flip its sign; multiply a step, scale its ΔH.
Bond enthalpies and enthalpy of formation
ΔH°_rxn ≈ Σ(bonds broken) − Σ(bonds formed). Bond enthalpies are averages, so results are approximate; ΔH°_f tables are exact per substance.
Quantitative skill-set
q = mcΔT; heating-curve segments; Hess cycles; ΔH°_f summation; bond-enthalpy estimates.
Exam traps
Standard state of an element: most stable form at 1 atm (graphite, O₂(g), Br₂(l)) with ΔH°_f = 0. Vaporization needs much more energy than fusion (breaking all IMFs). q = mcΔT applies only within a phase; at phase changes use ΔHfus/ΔHᵥap. Hess's law flips signs on reversal.