6.6 Introduction to Enthalpy of Reaction

The enthalpy change of a reaction, $\Delta H$, tells you how much heat a reaction releases or absorbs at constant pressure. A negative $\Delta H$ means heat leaves the system, so the reaction is exothermic; a positive $\Delta H$ means heat enters the system, so the reaction is endothermic. For AP Chemistry, match the sign of $\Delta H$ to the system's heat flow and the units in the balanced equation.

Enthalpy in AP Chemistry

In AP Chemistry, enthalpy of reaction is the heat absorbed or released by a chemical reaction at constant pressure. The sign of ΔH tells you the direction of heat flow: negative ΔH means the reaction is exothermic and releases heat to the surroundings, while positive ΔH means the reaction is endothermic and absorbs heat from the surroundings.

For AP Chem 6.6, the main calculation is q = n × ΔH. Use the balanced equation to understand what "per mole" refers to, convert the amount of reacting substance to moles, and keep track of whether your final heat value should be in J or kJ.

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Why This Matters for the AP Chemistry Exam

This topic gives you the tool to connect a balanced equation to a real heat measurement. On the AP Chemistry exam, you will apply mathematical routines to calculate or estimate the heat transferred and the overall enthalpy of a reaction, often by combining reaction data with the amount of substance in moles. Getting the sign of ΔH right and tracking units carefully (kJ vs kJ/mol) shows up in both calculation-style multiple-choice questions and free-response problems that ask you to support a claim with reasoning.

This topic also connects to calorimetry from earlier in Unit 6, since the heat you measure in an experiment links directly to the enthalpy change of a process. Later topics like bond enthalpies, enthalpy of formation, and Hess's Law all build on the idea that you can quantify ΔH.

Key Takeaways

• ΔH measures heat released or absorbed at constant pressure: negative means exothermic, positive means endothermic.
• Use q = n × ΔH to find total heat from the moles of reacting substance and the molar enthalpy of reaction.
• Watch your units: ΔH is usually in kJ/mol, while total heat q is in kJ or J.
• The chemical potential energy of products differs from reactants because bonds break and form, and that energy difference shows up as a temperature change.
• In an exothermic reaction, heat flows to the surroundings; in an endothermic reaction, heat flows from the surroundings, until thermal equilibrium is reached.
• You do not need to worry about the technical difference between enthalpy and internal energy. At constant pressure, ΔH equals the heat of reaction.

What Enthalpy of Reaction Means

Enthalpy (H) represents the heat content of a system at constant pressure. You cannot measure absolute enthalpy, but you can measure the enthalpy change (ΔH), which tells you how much heat energy is absorbed or released during a reaction at constant pressure.

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.

• In exothermic reactions, ΔH is negative. Heat is released by the system, and the products have lower enthalpy than the reactants.
• In endothermic reactions, ΔH is positive. Heat is absorbed by the system, and the products have higher enthalpy than the reactants.

These measurements assume the reaction occurs at constant pressure, which is the typical condition for reactions open to the atmosphere. You do not need the technical distinction between enthalpy and internal energy for the AP exam. At constant pressure, the enthalpy change equals the heat of reaction.

Exothermic and Endothermic Examples

• Exothermic: Combustion of methane: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g). ΔH is negative, so heat is released, often seen as a flame or rising temperature.
• Exothermic: Neutralization of an acid and base: HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l). ΔH is negative, and the solution warms up.
• Endothermic: Dissolution of ammonium nitrate: NH₄NO₃(s) → NH₄⁺(aq) + NO₃⁻(aq). ΔH is positive, so the solution gets colder as heat is absorbed.
• Endothermic: Synthesis of ammonia: N₂(g) + 3H₂(g) → 2NH₃(g). ΔH is positive.

Calculating Heat in Chemical Reactions

To calculate the heat (q) absorbed or released during a reaction, use:

q = n × ΔH

Where:

• q = heat absorbed or released (in J or kJ)
• n = number of moles of the reacting substance
• ΔH = molar enthalpy of reaction (in J/mol or kJ/mol)

Worked Example

For the combustion of methane: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g), ΔH = -890 kJ/mol

If you burn 2.5 mol of methane:

q = n × ΔH = 2.5 mol × (-890 kJ/mol) = -2225 kJ

The negative sign tells you that 2225 kJ of heat is released to the surroundings.

Thermal Equilibrium After Reactions

When the products of a reaction are at a different temperature than their surroundings, they exchange energy with the surroundings to reach thermal equilibrium.

• Exothermic reactions: Thermal energy is transferred to the surroundings as reactants convert to products. The products start warmer than the surroundings, then cool as heat flows out.
• Endothermic reactions: Thermal energy is transferred from the surroundings as reactants convert to products. The products start cooler than the surroundings, then warm as heat flows in.

Why the Temperature Changes

The chemical potential energy of the products differs from that of the reactants because of bond breaking and forming. That energy difference changes the kinetic energy of the particles, which you observe as a temperature change.

1. Bond breaking requires an energy input.
2. Bond forming releases energy.
3. The net difference between energy used to break bonds and energy released forming bonds determines ΔH.
4. That energy difference changes the kinetic energy of particles, which shows up as a temperature change.

So if more energy is released forming new bonds than is needed to break old bonds, the reaction is exothermic and the temperature rises. If breaking bonds takes more energy than forming bonds releases, the reaction is endothermic and the temperature drops.

How to Use This on the AP Chemistry Exam

Problem Solving

• Read the balanced equation and the given ΔH carefully. ΔH in kJ/mol is tied to the equation as written.
• Identify how many moles of the reacting substance you actually have, then apply q = n × ΔH.
• To solve for moles from a known heat, rearrange to n = q / ΔH.

Free Response

• When asked whether a reaction is exothermic or endothermic, state the sign of ΔH and connect it to heat flow and temperature change in the surroundings.
• Support claims with reasoning about bonds broken versus bonds formed when the question asks for an explanation, not just a value.
• Attend to significant figures and units in your final answer.

Common Trap

• Mixing up the sign of ΔH. Negative is exothermic (heat out), positive is endothermic (heat in).
• Forgetting that ΔH is per mole of reaction as written. If you double the amounts, you double the total heat.

Practice Problems

1. Calculate the heat released when 0.50 mol of propane combusts: C₃H₈(g) + 5O₂(g) → 3CO₂(g) + 4H₂O(g), ΔH = -2220 kJ/mol

   Solution: q = n × ΔH = 0.50 mol × (-2220 kJ/mol) = -1110 kJ

2. How many moles of NH₄NO₃ must dissolve to absorb 25 kJ of heat? NH₄NO₃(s) → NH₄⁺(aq) + NO₃⁻(aq), ΔH = +25.7 kJ/mol

   Solution: n = q / ΔH = 25 kJ / 25.7 kJ/mol = 0.97 mol

3. Identify whether each reaction is exothermic or endothermic:

   • 2H₂(g) + O₂(g) → 2H₂O(l), ΔH = -572 kJ/mol (Exothermic, ΔH < 0)
   • CaCO₃(s) → CaO(s) + CO₂(g), ΔH = +178 kJ/mol (Endothermic, ΔH > 0)

Key Terms Summary

• Heat (q): The transfer of thermal energy between a system and its surroundings, measured in joules (J) or kilojoules (kJ). Heat flows from higher temperature regions to lower temperature regions.
• Enthalpy (H): A quantity that represents the heat content of a system at constant pressure. You cannot measure absolute enthalpy, only enthalpy changes (ΔH).
• Enthalpy change (ΔH): The heat absorbed or released by a reaction at constant pressure. Negative means heat is released (exothermic), positive means heat is absorbed (endothermic).
• Molar enthalpy of reaction: The enthalpy change per mole of reaction as written in the balanced equation, usually in kJ/mol.

Common Misconceptions

• "Exothermic means the system gets colder." The system releases heat, so the surroundings warm up. In an exothermic reaction, the products start warmer than the surroundings.
• "ΔH is the same as temperature change." ΔH measures total energy flow at constant pressure, not temperature directly. Temperature change depends on how much substance is present and its specific heat.
• "A bigger amount of reactant changes ΔH." The molar enthalpy of reaction stays the same. What changes is the total heat q, because q = n × ΔH scales with moles.
• "Breaking bonds releases energy." Breaking bonds requires an energy input. Forming bonds releases energy. The net of both determines whether a reaction is exothermic or endothermic.
• "Positive ΔH means no heat is involved." Positive ΔH means heat is absorbed from the surroundings, which is just as real as heat being released.

Related AP Chemistry Guides

• Unit 6 Overview: Thermochemistry
• 6.1 Endothermic and Exothermic Processes
• 6.2 Energy Diagrams
• 6.4 Heat Capacity and Calorimetry
• 6.3 Heat Transfer and Thermal Equilibrium
• 6.7 Bond Enthalpies