Endothermic and Exothermic Processes

AP Chemistry· AP Chemistry CED — Thermodynamics· 14 min read

1. Core Definitions of Endothermic and Exothermic Processes★☆☆☆☆⏱ 3 min

Endothermic and exothermic processes describe the direction of heat flow between a thermodynamic system and its surroundings during any physical or chemical change. This is the core foundational concept for AP Chemistry Unit 6 Thermodynamics, and appears in both multiple-choice and free-response questions, often as a step in larger thermodynamics problems.

📘 Definition

Endothermic Process

A process where heat flows from the surroundings into the system (the reaction/change being studied). Applies to both chemical reactions and physical changes.

Example:

Photosynthesis, melting of ice, dissolution of ammonium nitrate in instant cold packs

📘 Definition

Exothermic Process

A process where heat flows out of the system into the surroundings. Applies to both chemical reactions and physical changes.

Example:

Combustion of fuel, condensation of water, freezing of liquid water to ice

2. Heat Flow and Enthalpy Sign Conventions★★☆☆☆⏱ 4 min

At constant pressure (the condition for nearly all reactions run in open laboratory containers), the enthalpy change of a process $Δ H$ equals the heat exchanged $q_{p}$ , so $Δ H = q_{p}$ . The universal chemistry convention is: heat gained by the system is positive, and heat lost by the system is negative.

This means for an endothermic process, where heat flows from surroundings into the system, $q > 0$ and $Δ H > 0$ . For an exothermic process, where heat flows from system to surroundings, $q < 0$ and $Δ H < 0$ . An intuitive check: if the reaction container feels warm to the touch, heat left the system to warm your hand (surroundings), so it is exothermic with negative $Δ H$ . If the container feels cold, heat left your hand (surroundings) to enter the system, so it is endothermic with positive $Δ H$ .

📐 Worked Example

A student activates an instant cold pack by mixing solid ammonium nitrate with room temperature water. The temperature of the resulting solution drops from 21°C to 3°C. Identify the dissolution process as endothermic or exothermic, and state the sign of $Δ H$ .

1
Define the system: the dissolution of ammonium nitrate in water is the system; the beaker, air, and the student's hand are the surroundings.
2
Track heat flow: the temperature of the solution (part of the system) dropped because thermal energy left the surroundings and entered the system to drive the dissolution process.
3
By convention, enthalpy change is measured relative to the system: heat gained by the system gives a positive $Δ H$ .
4
A process that absorbs heat from the surroundings is classified as endothermic. Final answer: Process is endothermic, $Δ H > 0$ .

Exam tip:

If an FRQ asks you to justify your classification, always explicitly mention heat flow relative to the system, not just the temperature change. AP graders require this reasoning for full credit.

3. Enthalpy Change from Bond Enthalpies★★★☆☆⏱ 5 min

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📘 Definition

Bond Enthalpy

BE

The average energy required to break one mole of a specific covalent bond in the gaseous state. Breaking any bond is always endothermic ( $Δ H > 0$ ), and forming any bond is always exothermic ( $Δ H < 0$ ).

Example:

The average bond enthalpy of an H-H single bond is 436 kJ/mol.

To calculate the total enthalpy change for a gas-phase reaction, we use the formula derived from summing bond breaking and bond formation enthalpies:

Delta H_{rxn} = \sum (\text{bond enthalpies of bonds broken}) - \sum (\text{bond enthalpies of bonds formed})

After calculating $Δ H_{r x n}$ , a positive value means the overall reaction is endothermic, and a negative value means the overall reaction is exothermic.

📐 Worked Example

Use the given average bond enthalpies to calculate $Δ H_{r x n}$ for the reaction $H_{2} (g) + Br_{2} (g) \to 2 HBr (g)$ , then classify the reaction. Bond enthalpies (kJ/mol): H-H = 436, Br-Br = 193, H-Br = 366.

1
Count bonds broken: 1 mol of H-H bonds and 1 mol of Br-Br bonds are broken. Sum of bonds broken =
2
$436 + 193 = 629 \text{ kJ}$
3
Count bonds formed: 2 mol of H-Br bonds are formed. Sum of bonds formed =
4
$2 \times 366 = 732 \text{ kJ}$
5
Calculate $Δ H_{r x n}$ using the formula:
6
$\Delta H_{rxn} = 629 - 732 = -103 \text{ kJ}$
7
Classify: $Δ H < 0$ , so the reaction is exothermic overall. Final answer: $Δ H = - 103$ kJ, reaction is exothermic.

Exam tip:

Always confirm the reaction is fully in the gas phase before using bond enthalpies. Bond enthalpy calculations for liquid or solid phase reactions omit phase change enthalpy terms, leading to incorrect results.

4. Energy Profile Diagrams and Process Classification★★☆☆☆⏱ 3 min

Energy profile (or reaction coordinate) diagrams plot the total enthalpy of the system versus reaction progress for a chemical process. These are regularly tested on the AP exam to assess understanding of the relationship between enthalpy change and process classification. The key feature for classification is the relative enthalpy of reactants versus products, following the definition:

Delta H = H_{final} - H_{initial} = H_{products} - H_{reactants}

For an exothermic process: products have lower enthalpy than reactants, so $Δ H < 0$ , and the diagram slopes downward from reactants to products. For an endothermic process: products have higher enthalpy than reactants, so $Δ H > 0$ , and the diagram slopes upward from reactants to products. Activation energy (the energy barrier from reactants to the activated complex) does not affect the overall classification of the reaction.

📐 Worked Example

A reaction coordinate diagram shows reactants at an enthalpy of 120 kJ/mol, an activated complex at 300 kJ/mol, and products at 75 kJ/mol. Calculate $Δ H_{r x n}$ , classify the reaction, and confirm your result.

1
Recall the definition of enthalpy change for a reaction:
2
$\Delta H_{rxn} = H_{products} - H_{reactants}$
3
Substitute the given values:
4
$\Delta H = 75 \text{ kJ/mol} - 120 \text{ kJ/mol} = -45 \text{ kJ/mol}$
5
A negative $Δ H$ means the system lost enthalpy overall, which corresponds to an exothermic reaction.
6
Confirm: products (75 kJ/mol) have lower enthalpy than reactants (120 kJ/mol), which matches the definition of an exothermic process. Final answer: $Δ H = - 45$ kJ/mol, reaction is exothermic.

Exam tip:

Do not confuse activation energy (the height of the reaction barrier) with the overall enthalpy change of the reaction. An exothermic reaction can still have a high activation energy (e.g., gasoline combustion requires a spark to start).

5. Common Pitfalls

Wrong move:

Classifying a cold pack reaction as exothermic because the solution temperature drops, concluding heat left the system.

Why:

Students confuse the temperature change of the system with the direction of heat flow. If the system's temperature drops, it absorbed heat from the surroundings.

Correct move:

Always state 'heat gained by system = endo, positive ΔH; heat lost by system = exo, negative ΔH' before classifying.

Wrong move:

Reversing the bond enthalpy formula, writing $Δ H = \sum B E_{f or m e d} - \sum B E_{b r o k e n}$ .

Why:

Students mix up that breaking bonds requires energy, so it adds a positive contribution to ΔH.

Correct move:

Use the mnemonic 'Broken First Minus Made' (B-F-M) to remember the formula order.

Wrong move:

Assigning a negative sign to the bond enthalpy of bonds broken.

Why:

Students incorrectly apply a negative sign to the endothermic bond breaking step, instead of using the formula that already accounts for sign.

Correct move:

Bond enthalpy values are always positive (they are the energy required to break bonds), so plug all BE values into the formula as positive.

Wrong move:

Classifying processes relative to the surroundings, getting the opposite sign: for example, saying an exothermic reaction has positive ΔH because the surroundings gain heat.

Why:

Students forget the AP Chem convention always uses the system as the reference point.

Correct move:

For any sign question, first write 'ΔH is relative to the system' in your working before assigning a sign.

Wrong move:

Calculating ΔH for an energy diagram as $H_{r e a c t an t s} - H_{p r o d u c t s}$ to force a negative sign for exothermic reactions.

Why:

Students rearrange the formula to get the sign they expect instead of following the state function rule.

Correct move:

Always use $Δ H = H_{f ina l} - H_{ini t ia l}$ , where final = products and initial = reactants, no exceptions.

6. Quick Reference Cheatsheet

Category	Formula / Rule	Notes
Endothermic process	$Δ H > 0$	Heat flows from surroundings into the system; surroundings cool. Always referenced to the system.
Exothermic process	$Δ H < 0$	Heat flows from system out to the surroundings; surroundings warm. Always referenced to the system.
Enthalpy change definition	$Δ H = H_{f ina l} - H_{ini t ia l}$	Applies to all enthalpy calculations; for reactions, final = products, initial = reactants.
Enthalpy from bond enthalpies	$Δ H_{r x n} = \sum B E_{b r o k e n} - \sum B E_{f or m e d}$	Only for gas-phase reactions. All bond enthalpy values are positive when plugged in.
Enthalpy and heat (constant P)	$Δ H = q_{p}$	Applies to all open-laboratory reactions at constant atmospheric pressure.
Endothermic energy profile	$H_{p r o d u c t s} > H_{r e a c t an t s}$	Diagram slopes upward from reactants to products; ΔH is positive.
Exothermic energy profile	$H_{p r o d u c t s} < H_{r e a c t an t s}$	Diagram slopes downward from reactants to products; ΔH is negative.

When this came up on past exams

AI-estimated based on syllabus patterns — cross-check with official past papers for accuracy. Use only as revision-focus signals.

2023 · MCQ
Classify dissolution, sign of ΔH
2022 · FRQ
Calculate ΔH from bond enthalpies

What's Next

Mastering endothermic and exothermic process classification and sign conventions is the non-negotiable foundation for all remaining thermodynamics topics in AP Chemistry Unit 6. Next you will apply these rules to calorimetry calculations, where you measure heat flow to determine enthalpy changes of reaction, then extend the concept to Hess's law for multi-step reactions. Incorrect sign conventions lead to wrong answers even with correct arithmetic, and this topic also underpins later concepts like enthalpy of formation, entropy, and Gibbs free energy where ΔH sign is critical for predicting reaction spontaneity.