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Unit 6 · Thermochemistry Unit Hub Flashcards Cheat Sheet Essentials Visual Review MC Practice SAQ Practice

AP Chemistry Unit 6 Visual Review

A topic-by-topic visual walkthrough of Thermochemistry — heat transfer, calorimetry, enthalpy, bond energies, and Hess's law.

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TOPIC 6.1 Endothermic & Exothermic Processes Exothermic (−ΔH) RELEASES energy to the surroundings. The system loses energy → ΔH is NEGATIVE. Products are LOWER in energy than reactants. Surroundings feel WARMER. Examples: combustion, freezing, most neutralizations, hand warmers. Energy comes out because forming bonds releases energy. Endothermic (+ΔH) ABSORBS energy from the surroundings. The system gains energy → ΔH is POSITIVE. Products are HIGHER in energy than reactants. Surroundings feel COOLER. Examples: melting, boiling, photosynthesis, instant cold packs. Energy goes in because breaking bonds requires energy. System vs. surroundings, and the sign convention ΔH is defined from the SYSTEM's point of view. Energy is always CONSERVED — it just moves between the system (the reaction) and the surroundings (everything else). If the system loses energy (−ΔH), the surroundings gain it. Overall energy of a reaction = (energy to break bonds) − (energy released forming bonds). A temperature change in the surroundings is the everyday evidence of an exo- or endothermic process. Exothermic releases energy (−ΔH, warmer); endothermic absorbs energy (+ΔH, cooler). The Review Hub · AP Chemistry Unit 6 TOPIC 6.2 Energy Diagrams Exothermic reactants products (lower) products below reactants · ΔH < 0 Endothermic reactants (lower) products products above reactants · ΔH > 0 Two things every energy diagram shows Activation energy (Eₐ) — the height from reactants up to the peak (transition state). It controls the RATE. Enthalpy change (ΔH) — the energy difference between products and reactants. It's exo- or endothermic. An energy diagram's shape shows both Eₐ (the hill) and ΔH (products vs. reactants). The Review Hub · AP Chemistry Unit 6 TOPIC 6.3 Heat Transfer & Thermal Equilibrium Heat flows hot → cold Heat (thermal energy) always transfers spontaneously from the HOTTER object to the COLDER one — never the reverse. It's driven by the difference in temperature (average KE). Temperature is intensive; heat (total energy) is extensive. Thermal equilibrium Heat keeps flowing until both objects reach the SAME temperature — thermal equilibrium. Then net heat flow stops. Same temperature ≠ same energy — depends on mass & material. Conservation of energy in heat transfer When a hot object meets a cold one in an insulated system, all the heat lost by the hot object equals the heat gained by the cold object: q_hot = − q_cold (energy conserved) The hot object cools (loses q, negative) as the cold object warms (gains q, positive). This is the principle behind calorimetry (Topic 6.4) — measuring heat by tracking a temperature change. The object with the larger heat capacity changes temperature less for the same amount of heat. Heat flows hot → cold until thermal equilibrium; heat lost = heat gained (energy conserved). The Review Hub · AP Chemistry Unit 6 TOPIC 6.4 Heat Capacity & Calorimetry q = m c ΔT What each term is q = heat (J) · m = mass (g) c = specific heat capacity (J/g·°C) ΔT = change in temperature (T_final − T_initial) Water's specific heat is high (4.18 J/g·°C) — it resists T change. Specific heat = resistance to T change A HIGH specific heat means a substance needs a lot of heat to change temperature (it warms/cools slowly). Calorimetry measures heat flow by tracking ΔT of a known mass of water in an insulated calorimeter. Worked example How much heat raises 50.0 g of water from 20.0 °C to 30.0 °C? q = m c ΔT = (50.0 g)(4.18 J/g·°C)(10.0 °C) = 2090 J (≈ 2.09 kJ) In a coffee-cup calorimeter, the heat released by a reaction equals the heat absorbed by the water: q_rxn = − q_water. A positive q means heat entered the substance; negative means heat left it. q = mcΔT — calorimetry finds a reaction's heat by measuring the temperature change of water. The Review Hub · AP Chemistry Unit 6 TOPIC 6.5 Energy of Phase Changes HEATING CURVE OF WATER melting (flat) boiling (flat) solid liquid gas temperature heat added → Flat plateaus = phase change During a phase change, temperature stays CONSTANT — the heat breaks IMFs instead of raising KE. On the sloped parts, q = mcΔT (temperature rises). Heat of fusion & vaporization ΔH_fus — energy to melt (solid → liquid) ΔH_vap — energy to boil (liquid → gas) ΔH_vap > ΔH_fus (all IMFs break in boiling). q = n × ΔH for the phase-change plateaus. Stronger IMFs → higher melting/boiling points AND larger heats of fusion/vaporization. During a phase change, heat breaks IMFs and temperature stays constant (the flat plateaus). The Review Hub · AP Chemistry Unit 6 TOPIC 6.6 Enthalpy of Reaction (ΔH_rxn) Enthalpy (H) is heat at constant pressure ΔH_rxn is the heat absorbed or released by a reaction at constant pressure. It's a STATE FUNCTION — its value depends only on the initial and final states, NOT on the path taken (the key idea behind Hess's law, Topic 6.9). Negative ΔH = exothermic; positive ΔH = endothermic. ΔH scales with amount ΔH is proportional to the moles reacted. Double the reaction → double the ΔH. Report it per the balanced equation (kJ/mol_rxn). This lets you use ΔH like a stoichiometric quantity. Reversing a reaction flips the sign If a forward reaction is exothermic (ΔH = −x kJ), the reverse reaction is endothermic with ΔH = +x kJ. Example: combustion of CH₄ has ΔH = −890 kJ/mol. These two rules are what make Hess's law work. There are three main ways to find ΔH_rxn: calorimetry, bond enthalpies (6.7), and Hess's law / formation enthalpies (6.8–6.9). ΔH_rxn is a state function — path-independent, scales with moles, and flips sign when reversed. The Review Hub · AP Chemistry Unit 6 TOPIC 6.7 Bond Enthalpies ΔH = Σ(bonds broken) − Σ(bonds formed) Breaking absorbs, forming releases Breaking bonds REQUIRES energy (endothermic, +). Forming bonds RELEASES energy (exothermic, −). ΔH_rxn is the net difference between the two. Bond enthalpies are always positive (energy to break the bond). How to read the sign If the products' bonds are STRONGER (release more than was spent breaking reactants), ΔH is negative → exothermic. Bond enthalpies give an ESTIMATE — they use average values, so they're less exact than formation enthalpies. Worked example — H₂ + Cl₂ → 2 HCl Bonds broken: 1 H–H (436) + 1 Cl–Cl (243) = 679 kJ absorbed Bonds formed: 2 H–Cl (2 × 431) = 862 kJ released ΔH = 679 − 862 = −183 kJ (exothermic — the HCl bonds are stronger) Only count bonds that actually break/form; unchanged bonds cancel out. ΔH = bonds broken − bonds formed — breaking costs energy, forming releases it. The Review Hub · AP Chemistry Unit 6 TOPIC 6.8 Enthalpy of Formation ΔH_rxn = Σ ΔH°f(products) − Σ ΔH°f(reactants) Standard enthalpy of formation (ΔH°f) The heat change when 1 mole of a compound forms from its ELEMENTS in their standard states. ΔH°f of any pure element = 0. e.g., O₂(g), N₂(g), C(graphite), Fe(s) all have ΔH°f = 0. "Products minus reactants" Look up each substance's ΔH°f in a table, multiply by its coefficient, then subtract reactants from products. This gives a very accurate ΔH_rxn (unlike bond enthalpies, which use averages). Worked example — CH₄ + 2 O₂ → CO₂ + 2 H₂O ΔH°f: CH₄ = −75, O₂ = 0, CO₂ = −394, H₂O = −286 (kJ/mol) Products: (−394) + 2(−286) = −966 · Reactants: (−75) + 2(0) = −75 ΔH_rxn = −966 − (−75) = −891 kJ (exothermic combustion) Don't forget the coefficients — each ΔH°f is multiplied by the number of moles in the equation. ΔH_rxn = Σ ΔH°f products − Σ ΔH°f reactants; elements in standard state have ΔH°f = 0. The Review Hub · AP Chemistry Unit 6 TOPIC 6.9 Hess's Law ΔH is path-independent, so steps add up Because enthalpy is a state function, the ΔH of an overall reaction equals the SUM of the ΔH values of the steps used to get there — no matter what route you take. This lets you find a hard-to-measure ΔH from known ones. Whether a reaction happens in one step or ten, the total energy change is the same. Two manipulation rules 1. REVERSE a step → flip the sign of its ΔH. 2. MULTIPLY a step by a factor → multiply its ΔH too. Then add the adjusted steps so intermediates cancel out. Target reaction on top; arrange known steps to match it. The strategy Line up the given equations so that: • each reactant/product ends up on the correct side • intermediates appear on both sides and CANCEL Add the ΔH values of your adjusted steps → target ΔH. The formation-enthalpy method (Topic 6.8) is really just a shortcut application of Hess's law. Hess's law: add up known steps' ΔH (flip/scale as needed) to find an overall ΔH. The Review Hub · AP Chemistry Unit 6
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This is great for review the night before the exam — fast, visual, and covers everything you need to remember about Unit 6's thermochemistry content.