6.6 Introduction to Enthalpy of Reaction

Enthalpy of reaction (ΔHrxn) is the heat change for a chemical reaction at constant pressure—negative ΔH means heat released (exothermic), positive means heat absorbed (endothermic). It comes from differences in chemical potential energy when bonds break and form; that energy shows up as temperature change of the products and surroundings (CED 6.6.A.1–6.6.A.3). For AP problems you often use molar enthalpy: q = n × ΔHrxn (or measure q with a coffee-cup calorimeter and relate it to ΔH at constant pressure). Knowing ΔHrxn lets you predict whether a reaction warms or cools its surroundings, calculate heat for a given amount of reactant, and connect bond energetics to macroscopic temperature changes—all skills tested in Unit 6 thermochemistry (practice and calorimetry questions appear on the exam). For a focused review check the Topic 6.6 study guide (https://library.fiveable.me/ap-chemistry/unit-6/intro-enthalpy-reaction/study-guide/MZToO4yTqx0BDiUxmvhN), the Unit 6 overview (https://library.fiveable.me/ap-chemistry/unit-6), and thousands of practice problems (https://library.fiveable.me/practice/ap-chemistry).

You calculate heat (q) by linking the reaction’s molar enthalpy (ΔHrxn) to how many moles actually react and using sign conventions from the CED. Steps: 1) Get ΔHrxn (kJ per mole of reaction). Remember at constant pressure qp = ΔH (CED 6.6.A.1). Negative ΔH = exothermic (releases heat); positive = endothermic (absorbs heat). 2) Use stoichiometry to find moles of the limiting reactant that react (n). 3) Multiply: q = n × ΔHrxn (same units). Example: if ΔHrxn = −286 kJ/mol and 2.0 mol react, q = 2.0 × (−286) = −572 kJ released. If you measured temperature change in a coffee-cup calorimeter, use qreaction = −qsolution and qsolution = m·s·ΔT (mass × specific heat × ΔT) to find qreaction (CED 6.6.A.2, calorimetry keywords). For more examples and practice problems, check the Topic 6.6 study guide (https://library.fiveable.me/ap-chemistry/unit-6/intro-enthalpy-reaction/study-guide/MZToO4yTqx0BDiUxmvhN), the Unit 6 overview (https://library.fiveable.me/ap-chemistry/unit-6), or lots of practice questions (https://library.fiveable.me/practice/ap-chemistry).

Exothermic vs. endothermic is all about ΔH (enthalpy change) and heat flow at constant pressure. If ΔH is negative, the reaction is exothermic: chemical potential energy drops as bonds form more strongly than they were broken, thermal energy is released to the surroundings, and the surroundings’ temperature usually rises. If ΔH is positive, the reaction is endothermic: the system absorbs thermal energy from the surroundings (breaking bonds costs more energy than forming), so the surroundings cool and the products end up with higher chemical potential energy. On the AP exam you can treat ΔHrxn as the heat at constant pressure (qp = ΔH). For calculations, use q = n·ΔHrxn or calorimetry (q = m·c·ΔT) to relate moles and temperature changes. For a quick refresher, see the Topic 6.6 study guide (https://library.fiveable.me/ap-chemistry/unit-6/intro-enthalpy-reaction/study-guide/MZToO4yTqx0BDiUxmvhN) and more Unit 6 resources (https://library.fiveable.me/ap-chemistry/unit-6); practice problems are at (https://library.fiveable.me/practice/ap-chemistry).

Temperature changes in reactions because chemical energy is redistributed as heat. At the AP level you use enthalpy change (ΔH) to describe that heat at constant pressure: ΔH < 0 (exothermic) means the reaction releases thermal energy to the surroundings and their temperature rises; ΔH > 0 (endothermic) means the reaction absorbs thermal energy and surroundings cool (CED 6.6.A.1–2). Microscopically, breaking bonds requires energy and forming bonds releases energy, so the net difference changes the kinetic energy of molecules—that change in kinetic energy shows up as a temperature change (CED 6.6.A.3). After the reaction the system and surroundings exchange heat until thermal equilibrium is reached (CED 6.6.A.2). For practice on calculations using q = n·ΔHrxn and calorimetry, see the Topic 6.6 study guide (https://library.fiveable.me/ap-chemistry/unit-6/intro-enthalpy-reaction/study-guide/MZToO4yTqx0BDiUxmvhN) and >1,000 practice problems (https://library.fiveable.me/practice/ap-chemistry).

Negative enthalpy (ΔH < 0) means heat is released—that’s an exothermic reaction. Positive enthalpy (ΔH > 0) means heat is absorbed—that’s endothermic. For most AP Chem work (constant pressure), ΔHrxn = qp, so the sign of ΔH tells you the direction of heat flow with the surroundings (negative → heat to surroundings, positive → heat from surroundings). Example from the CED: ΔH° = −198 kJ/mol means 198 kJ of heat is released per mole-reaction. When calculating heat for a given amount, use q = n·ΔH(molar) (watch stoichiometry and units). This sign convention and using ΔH at constant pressure are what the AP exam expects you to know (Topic 6.6). For a quick refresher, see the Topic 6.6 study guide (https://library.fiveable.me/ap-chemistry/unit-6/intro-enthalpy-reaction/study-guide/MZToO4yTqx0BDiUxmvhN) and more practice at (https://library.fiveable.me/practice/ap-chemistry).

Think of molar enthalpy (ΔHrxn) as a conversion factor between moles and heat. Steps: 1. Start with the molar enthalpy given (often in kJ per mol of reaction). Use the AP sign convention: negative ΔH = exothermic (heat released), positive = endothermic (absorbed). 2. Convert ΔH to joules if you need J: 1 kJ = 1000 J. 3. Use q = n · ΔHrxn to get heat (q). If ΔH was in kJ/mol and you want J, do q(J) = n(mol) · ΔH(kJ/mol) · 1000 J/kJ. 4. If you start with mass, convert mass → moles via n = mass / molar mass, then use step 3. 5. In calorimetry problems, you may also set q(reaction) = −q(solution) and compute q(solution) = m · c · ΔT (m in g, c in J/g·°C). AP tip: the CED expects you to calculate q from moles and molar enthalpy (Learning Objective 6.6.A). For extra practice, see the Topic 6.6 study guide (https://library.fiveable.me/ap-chemistry/unit-6/intro-enthalpy-reaction/study-guide/MZToO4yTqx0BDiUxmvhN) and thousands of practice questions (https://library.fiveable.me/practice/ap-chemistry).

When a reaction happens, energy moves because chemical bonds change. Breaking bonds requires energy input (that’s endothermic) and forming new bonds releases energy (that’s exothermic). At constant pressure the heat exchanged is the enthalpy change ΔH: if the energy released by bond formation > energy absorbed to break bonds, ΔH is negative and the reaction is exothermic (surroundings heat up). If bond breaking costs more than formation releases, ΔH is positive and the reaction is endothermic (system absorbs heat from the surroundings, so temperature drops). AP-style rule of thumb: approximate ΔHrxn ≈ (energy of bonds broken) − (energy of bonds formed). The difference shows up as a change in kinetic energy of particles and thus temperature change (thermal equilibrium is reached afterward). For review, see the Topic 6.6 study guide (https://library.fiveable.me/ap-chemistry/unit-6/intro-enthalpy-reaction/study-guide/MZToO4yTqx0BDiUxmvhN) and more Unit 6 resources (https://library.fiveable.me/ap-chemistry/unit-6). Practice problems are at (https://library.fiveable.me/practice/ap-chemistry).

They exchange energy because heat flows from the hotter place to the colder place until both have the same temperature—that’s thermal equilibrium. Temperature measures average kinetic energy of particles; a temperature difference means particles on one side have higher average kinetic energy, so collisions transfer energy to the cooler side. As energy flows, kinetic energies change until the averages equalize (no net heat flow). In chemical reactions this shows up as enthalpy change (ΔH): at constant pressure, the heat released or absorbed by the system goes into or comes from the surroundings (exothermic → surroundings gain thermal energy; endothermic → surroundings lose thermal energy). The driving “why” is statistical/thermodynamic: spontaneous heat flow increases total entropy of the universe. This idea ties directly to Topic 6.6 (ΔH and heat flow) and Topic 6.3 (thermal equilibrium). For a clear AP-aligned review, see the Topic 6.6 study guide (https://library.fiveable.me/ap-chemistry/unit-6/intro-enthalpy-reaction/study-guide/MZToO4yTqx0BDiUxmvhN) and more Unit 6 resources (https://library.fiveable.me/ap-chemistry/unit-6).

Look at the sign of the reaction enthalpy ΔH: negative ΔH = exothermic (releases heat to the surroundings); positive ΔH = endothermic (absorbs heat from the surroundings). At AP level you’ll see this used a few ways: - If ΔHrxn is given (or ΔH°f for products/reactants), sign tells you directly. Use Hess’s law or ΔH°rxn = ΣΔH°f(products) − ΣΔH°f(reactants) to find it. - If you only have bond changes, remember: breaking bonds costs energy (endothermic), forming bonds releases energy (exothermic). If total energy released by bond formation > energy required to break bonds → net release (exothermic). - Common patterns: most combustion and formation reactions are exothermic; dissolving salts or melting can be endothermic or exothermic depending on enthalpy terms—don’t assume without data. AP tip: AP problems expect you to state the sign of ΔH and connect it to heat flow at constant pressure (qp = ΔH) and to bond-making/breaking (CED 6.6.A). For a quick refresher, see the Topic 6.6 study guide (https://library.fiveable.me/ap-chemistry/unit-6/intro-enthalpy-reaction/study-guide/MZToO4yTqx0BDiUxmvhN) and try practice problems (https://library.fiveable.me/practice/ap-chemistry).

Chemical potential energy is the stored energy in bonds of reactants/products; kinetic energy is the motion of particles (what we feel as temperature). During a reaction, bonds break and form so the chemical potential energy of products differs from reactants (CED 6.6.A.3). The energy difference shows up as a change in kinetic energy of molecules or as heat exchanged with the surroundings: - Exothermic: products have lower chemical potential energy → excess energy becomes kinetic energy of particles or heat released (ΔH < 0), so temperature of surroundings rises. - Endothermic: products have higher chemical potential energy → kinetic energy is used to make that potential energy (ΔH > 0), so surroundings cool. At constant pressure, the enthalpy change ΔH equals the heat q absorbed or released (CED 6.6.A.1–2). For more examples and practice, see the Topic 6.6 study guide (https://library.fiveable.me/ap-chemistry/unit-6/intro-enthalpy-reaction/study-guide/MZToO4yTqx0BDiUxmvhN) and Unit 6 overview (https://library.fiveable.me/ap-chemistry/unit-6). For lots of practice problems, check (https://library.fiveable.me/practice/ap-chemistry).

Because most reactions you study in AP Chem happen open to the atmosphere (like in beakers or coffee-cup calorimeters), pressure stays basically constant. Enthalpy H is defined so that the heat exchanged at constant pressure qp equals the change in enthalpy ΔH. Mathematically, ΔH = ΔU + Δ(PV); if P is constant, Δ(PV) simplifies and ΔH = q_p, so the heat measured is the enthalpy change. That’s why the CED defines ΔHrxn as the heat at constant pressure (6.6.A.1) and why calorimetry problems use constant-pressure setups. The AP exam won’t test the technical differences between ΔH and ΔU (exclusion statement), so focus on using ΔH = q_p for typical lab problems. For a quick refresher, see the Topic 6.6 study guide (https://library.fiveable.me/ap-chemistry/unit-6/intro-enthalpy-reaction/study-guide/MZToO4yTqx0BDiUxmvhN) or the Unit 6 overview (https://library.fiveable.me/ap-chemistry/unit-6). For more practice, check Fiveable’s problem set library (https://library.fiveable.me/practice/ap-chemistry).

Think of bonds as stored chemical potential energy. Breaking bonds requires energy input (you’re raising potential energy), so the system absorbs heat from its surroundings → endothermic (ΔH > 0), surroundings cool. Forming bonds releases energy (potential energy drops), so energy is dumped into particle motion → exothermic (ΔH < 0), surroundings warm. At constant pressure the enthalpy change ΔHrxn equals the heat q exchanged (CED 6.6.A.1), and any ΔH difference between reactants and products becomes kinetic energy of particles, which shows up as a temperature change until thermal equilibrium is reached (CED 6.6.A.2–3). So: breaking = energy in (cooling); forming = energy out (heating). If more energy is released forming bonds than used breaking them, net heat is released (exothermic). For more practice and AP-aligned explanations, see the Topic 6.6 study guide (https://library.fiveable.me/ap-chemistry/unit-6/intro-enthalpy-reaction/study-guide/MZToO4yTqx0BDiUxmvhN) and try practice problems (https://library.fiveable.me/practice/ap-chemistry).

Quick recipe for AP-style enthalpy problems (Topic 6.6): 1. Identify what you're given: a molar enthalpy ΔHrxn (kJ per mol rxn), a mass or moles of a reactant, or calorimetry data (m, c, ΔT). Note AP learning goal 6.6.A: relate q to moles and molar enthalpy. 2. If you have moles of a substance: use stoichiometry. q = n (mol of reaction or of a species) × ΔHrxn (kJ/mol). Watch whether ΔH is “per mole of reaction” or per mole of a reactant—scale by the balanced equation. 3. If you have mass: convert mass → moles using molar mass, then do step 2. 4. If it’s calorimetry (coffee-cup, constant pressure): qsys = −qsurr. Use q = m c ΔT for the solution/spoon/etc. At constant pressure qp = ΔH, so you can find ΔHrxn by dividing q (kJ) by moles of limiting reactant. 5. Sign and units: negative q or ΔH = exothermic (heat released); positive = endothermic (absorbed). Keep units consistent (J ↔ kJ). For worked examples and practice aligned to the CED, check the Topic 6.6 study guide (https://library.fiveable.me/ap-chemistry/unit-6/intro-enthalpy-reaction/study-guide/MZToO4yTqx0BDiUxmvhN), the whole Unit 6 overview (https://library.fiveable.me/ap-chemistry/unit-6), and drill lots of problems at (https://library.fiveable.me/practice/ap-chemistry).

Molar enthalpy (ΔHrxn) tells you how much heat is released or absorbed per mole of reaction as written at constant pressure. To get the actual heat q exchanged, multiply ΔHrxn by the number of moles (n) that reacted and by any stoichiometric factor: q = n × ΔHrxn. Sign matters: ΔH < 0 → q is negative (exothermic, heat flows to surroundings); ΔH > 0 → q is positive (endothermic, heat flows from surroundings). Example: if ΔH = −100 kJ per mole of reaction and 2.0 mol react, q = 2.0 × (−100 kJ/mol) = −200 kJ released. In calorimetry (e.g., a coffee-cup calorimeter) you measure the temperature change of the surroundings to find q and relate it back to moles via that equation (LO 6.6.A in the CED). For practice applying this with stoichiometry and calorimetry problems, see the Topic 6.6 study guide (https://library.fiveable.me/ap-chemistry/unit-6/intro-enthalpy-reaction/study-guide/MZToO4yTqx0BDiUxmvhN) and more problems at (https://library.fiveable.me/practice/ap-chemistry).

Exothermic vs. endothermic comes down to energy bookkeeping for bonds and thermal equilibrium. If forming bonds in the products releases more chemical potential energy than was needed to break reactant bonds, ΔH is negative—that excess energy appears as thermal energy, so the products end up hotter and transfer heat to the surroundings until thermal equilibrium is reached. In an endothermic reaction, breaking bonds (and raising product potential energy) requires more energy than bond formation releases, so ΔH is positive—the system absorbs thermal energy from the surroundings, cooling them as it moves toward equilibrium. At constant pressure, ΔH = qp, so the sign of ΔH tells you whether heat leaves or enters the system (CED 6.6.A.1–6.6.A.3). For a quick refresher, see the Topic 6.6 study guide (https://library.fiveable.me/ap-chemistry/unit-6/intro-enthalpy-reaction/study-guide/MZToO4yTqx0BDiUxmvhN) and Unit 6 resources (https://library.fiveable.me/ap-chemistry/unit-6). Need practice problems? Try the AP problem set (https://library.fiveable.me/practice/ap-chemistry).