AP Chemistry Endothermic vs Exothermic Reactions

Four ways to classify a process, with a diagram, a bond enthalpy calculation, and practice questions.

The terms are easy to define and easy to flip in your head, because the sign of ΔH is measured from the system and the thermometer reads the surroundings. The way to stay accurate is to have more than one route to the same answer and check them against each other. This page gives four routes, then uses them on a cold pack, an energy diagram, a reaction estimated from bond enthalpies, and a phase change. The Unit 6 review covers where this sits in the unit.

Four Ways to Tell an Endothermic Process From an Exothermic One

The system is the reaction or process itself and the surroundings are everything else, such as the water in a calorimeter. Energy that leaves one enters the other, so the two always have opposite signs: if the surroundings gain heat, the system's q is negative. That is the same sign logic used in calorimetry.

Worked Example 1: Classifying a Cold Pack From Its Temperature Change

When ammonium nitrate dissolves in 100.0 g of water in a coffee-cup calorimeter, the water temperature falls by 3.5°C. Is dissolving endothermic or exothermic, and what is q for the dissolving?

The surroundings (the water) got colder, so they lost energy and the process absorbed it: endothermic.
q_water = mcΔT = (100.0 g)(4.18 J/(g·°C))(−3.5°C) = −1463 J, about −1.5 kJ.
The system's heat is the opposite: q_dissolving = +1.5 kJ. A positive value means the process gained energy, which agrees with ΔH > 0 at constant pressure.

Worked Example 2: Reading ΔH From an Energy Diagram

Each diagram below shows reactant and product energies. Find ΔH and classify each reaction.

ΔH = energy of products − energy of reactants.
Left: 30 − 70 = −40 kJ/mol, exothermic.
Right: 70 − 30 = +40 kJ/mol, endothermic.

The peak between the levels is the activation energy, which sets the speed and not the sign. Both reactions here have a barrier, and an exothermic reaction can still need a push to start. The barrier is the subject of activation energy.

Worked Example 3: Estimating ΔH From Bond Enthalpies

Estimate ΔH for H₂(g) + Cl₂(g) → 2 HCl(g) using bond enthalpies of 436 kJ/mol for H–H, 242 kJ/mol for Cl–Cl, and 431 kJ/mol for H–Cl.

Bonds broken (energy absorbed): 1 H–H + 1 Cl–Cl = 436 + 242 = 678 kJ.
Bonds formed (energy released): 2 H–Cl = 2 × 431 = 862 kJ.
ΔH = bonds broken − bonds formed = 678 − 862 = −184 kJ, exothermic.

More energy is released forming two H–Cl bonds than was spent breaking H–H and Cl–Cl. Table values differ slightly from book to book, and the exam supplies the ones a question needs. If you reverse the reaction, 2 HCl → H₂ + Cl₂, the magnitude is the same and the sign flips to +184 kJ, which is how Hess's law treats a reverse step. The result is an estimate because bond enthalpies are averages for gas-phase molecules.

Worked Example 4: Explaining a Phase Change by the Strength of Interactions

Is the condensation of water vapor to liquid water endothermic or exothermic? Explain in terms of the particles.

In the vapor the molecules are far apart with almost no attraction between them. When they condense, hydrogen bonds form between neighboring molecules. Forming attractions releases energy, so the particles end up at lower energy and the surroundings are warmed: exothermic, ΔH < 0. The reverse change, vaporization, must pull those interactions apart, so it is endothermic.

This is the strength-of-interactions argument the course uses for physical changes, and the same bookkeeping explains why temperature stays flat during a plateau on a heating curve.

Five Endothermic and Exothermic Practice Questions With Answers

1. A reaction has ΔH = −92 kJ. Classify it and say what happens to the temperature of the surroundings.
   Show answer

   Exothermic. Energy is released, so the surroundings warm up.
2. Reactants are at 120 kJ/mol and products at 75 kJ/mol. Find ΔH.
   Show answer

   ΔH = 75 − 120 = −45 kJ/mol, so the reaction is exothermic.
3. A beaker feels cold after two solids are mixed. Is the process endothermic or exothermic, and what is the sign of q for the process?
   Show answer

   Endothermic. The surroundings lost energy, so q for the process is positive.
4. Estimate ΔH for H₂ + F₂ → 2 HF using H–H 436, F–F 155, and H–F 567 kJ/mol.
   Show answer

   Broken: 436 + 155 = 591 kJ. Formed: 2 × 567 = 1134 kJ. ΔH = 591 − 1134 = −543 kJ, exothermic.
5. Is melting endothermic or exothermic? Explain.
   Show answer

   Endothermic (ΔH > 0). Energy is absorbed to overcome the attractions holding the particles in the solid.

Common Endothermic and Exothermic Mistakes

• Reading the sign from the surroundings instead of the system. A cold beaker means the process is endothermic, with ΔH > 0.
• Subtracting bond energies the wrong way. For ΔH, use bonds broken minus bonds formed; reversing it flips the sign.
• Assuming exothermic means fast or spontaneous. The sign of ΔH says nothing about speed and does not settle spontaneity.
• Calling breaking bonds exothermic. Breaking bonds always absorbs energy; forming bonds always releases it.
• Forgetting to multiply bond enthalpies by the coefficient. Two H–Cl bonds are formed in the example, not one.

Once classification is automatic, the enthalpy calculations in Hess's law and calorimetry follow from the same sign convention. To see how your unit-by-unit practice might convert into an overall result, the AP Chem Score Calculator is a quick way to find out.

Related Resources

• AP Chem Score Calculator
• Unit 6 Review: Thermochemistry
• AP Chemistry Calorimetry and Specific Heat
• AP Chemistry Activation Energy
• AP Chemistry Hess's Law
• AP Chemistry Heating Curve
• AP Chemistry Study Guide

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