Introduction to Enthalpy of Reaction
Scale reaction heat with the balanced equation and distinguish reaction energy from the surroundings’ warming.
AP Chemistry Topic 6.6 · CED printed p. 118 / PDF p. 125. This teaching, practice and feedback were added for the AP edition; original author material remains separately traceable.
Before you start
- System and surroundings as a chosen boundary
- Moles, balanced equations and limiting reactants
- Intermolecular versus intramolecular interactions
- Activation energy versus overall energy change
At constant pressure, the enthalpy change gives the heat absorbed or released for the reaction amount represented by the balanced equation, under the stated conditions. For 2H2(g) + O2(g) → 2H2O(l), ΔH =−571.6 kJ means two moles H2 reacting transfer −571.6 kJ; it is not −571.6 kJ per mole H2. Use reaction extent: q =ξΔH, where ξ equals reacting moles divided by their coefficient. Scale every coefficient and ΔH together.
An exothermic reaction creates products with lower chemical potential energy than reactants. Some of that difference can appear as increased particle thermal motion before heat transfers to cooler surroundings. An endothermic reaction draws energy from surroundings, which may cool. Heat exchange continues until thermal equilibrium is reached. In calorimetry, the measured solution response is used to infer reaction heat; it is not a second independent source of energy.
Bond breaking requires energy input and bond formation releases energy. The net change depends on both, not on breaking alone. At the same final temperature reference, different bonded configurations have different chemical potential energies; during the actual event those differences can shift kinetic energy and temperature. The AP course uses constant-pressure reaction heat as ΔH and does not assess the technical distinction between enthalpy and internal energy here. Conditions, phase labels and the amount reacting must still be stated.
Let's work through it
For 2H2(g) + O2(g) → 2H2O(l), ΔH =−571.6 kJ. A supplied reaction consumes 0.150 mol H2 with sufficient oxygen.
- Reaction extent ξ =0.150/2 =0.0750 mol of reaction as written.
- q =0.0750(−571.6) =−42.9 kJ. Surroundings gain +42.9 kJ in an ideal heat-only balance.
- Writing H2 +½O2 → H2O(l) halves ΔH to −285.8 kJ; using 0.150 mol with this form gives the same q.
A reaction’s new bond formation releases more energy than the initial bond breaking consumes.
- The final chemical configuration has lower potential energy than the initial one.
- Before equilibration, some released energy can increase thermal motion and heat the mixture.
- When the mixture is hotter than its surroundings, energy transfers outward until temperatures equalize. The reaction is exothermic even though the measured mixture initially warms.
Try it from memory
Pause & retrieve
Try it before you peek.
Write your answer and the reason from memory. Then compare with the worked answer. Your self-assessment is a study decision, not an automatically graded score.
Question 1
If an equation’s coefficients are doubled, what happens to its ΔH?
See every recall question with its answer
These references preserve each prompt and its answer for direct evidence links. Try the interactive questions above before opening the worked reasoning.
Retrieval 1
If an equation’s coefficients are doubled, what happens to its ΔH?
Worked answer and reason
ΔH doubles with the amount represented.
Enthalpy transfer scales with the reaction amount.
Retrieval 2
Why can bond breaking not be the reason an exothermic reaction releases heat?
Worked answer and reason
Bond breaking requires energy; the larger release from forming new interactions outweighs that input.
Net heat reflects competing breaking and forming contributions.
Catch the tempting mistake
- ΔH printed next to an equation is automatically per mole of any reactant.
- It refers to the stoichiometric reaction as written; divide reacting moles by the relevant coefficient.
- Breaking bonds releases stored energy.
- Separating bonded atoms requires input; formation releases energy.
Put it to work
Give the chemical reason as well as the calculation. Worked solutions are on a separate page.
- Unit 6 problem 9 · ap application · multiple
- Unit 6 problem 10 · ap application · data
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For the teacher
- Ask students to name the mole basis beside each ΔH value.
- Require phase labels because gas and liquid products can have different enthalpies.
- Use q009’s separate post-reaction time sequence to check eventual temperature equality. Ask students to distinguish finite surroundings from a constant-temperature reservoir and thermal equilibrium from reaction equilibrium.
Source objectives and scope
Source edition: ap-chemistry-ced-fall2024-local-b5dfe867. See this topic in the course outline →
- 6.6.A
- Calculate the heat q absorbed or released by a system undergoing a chemical reaction in relationship to the amount of the reacting substance in moles and the molar enthalpy of reaction.
- 6.6.A.1
- The enthalpy change of a reaction gives the amount of heat energy released (for negative values) or absorbed (for positive values) by a chemical reaction at constant pressure.
- 6.6.A.2
- When the products of a reaction are at a different temperature than their surroundings, they exchange energy with the surroundings to reach thermal equilibrium. Thermal energy is transferred to the surroundings as the reactants convert to products in an exothermic reaction. Thermal energy is transferred from the surroundings as the reactants convert to products in an endothermic reaction.
- 6.6.A.3
- The chemical potential energy of the products of a reaction is different from that of the reactants because of the breaking and forming of bonds. The energy difference results in a change in the kinetic energy of the particles, which manifests as a temperature change.
- The technical distinctions between enthalpy and internal energy will not be assessed on the AP Exam. Most reactions studied at the AP level are carried out at constant pressure, where the enthalpy change of the process is equal to the heat (and by extension, the energy) of reaction.
Science references for this teaching
- College Board AP Chemistry Course and Exam Description, Fall 2024 — Exact unit learning requirements and exclusions; supplied PDF is pinned by the integration catalog.
- OpenStax Chemistry 2e: Enthalpy — Scientific check: stoichiometric enthalpy, formation values, sign reversal/scaling and Hess cycles.
Glossary and deeper reading · 5 refreshers
These editorial refreshers supplement the reading. Definitions and examples are added guidance; original source wording and recorded answers remain in their original context.
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Mole
A mole is an amount of substance containing exactly 6.02214076 × 10²³ specified entities. The entities can be atoms, molecules, ions, or formula units, so the counted entity must be named. For example, one mole of oxygen molecules contains two moles of oxygen atoms.
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Limiting reactant
The limiting reactant is consumed first under the specified complete-reaction model and limits the product amount. Identifying it requires comparing available amounts with the balanced equation's ratios. For example, abundant hydrogen cannot produce more water after the available oxygen has been consumed.
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Activation energy
Activation energy is the energy barrier associated with a reaction pathway in an introductory energy-profile model. It concerns the pathway and is distinct from the reactant-to-product energy change. For example, a thermodynamically favored reaction can remain slow when its activation barrier is large.
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Exothermic process
An exothermic process transfers heat from the chosen system to its surroundings. Under constant-pressure conditions with only pressure–volume work, its enthalpy change is negative. For example, freezing water releases heat to the surroundings.
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Endothermic process
An endothermic process absorbs heat into the chosen system. Under constant-pressure conditions with only pressure–volume work, its enthalpy change is positive. For example, melting ice absorbs heat while the phase change occurs.