AP Chemistry: Endothermic vs Exothermic, Enthalpy (ΔH)

AP Chemistry – Endothermic vs Exothermic, Enthalpy (ΔH)

AP Chemistry Study Guide: Endothermic vs. Exothermic & Enthalpy (ΔH)

What This Is

Endothermic and exothermic reactions describe whether a process absorbs or releases heat energy. Enthalpy (ΔH) is the heat content of a system at constant pressure—it tells us how much energy is gained or lost in a reaction. This is critical for the AP exam because it appears in thermochemistry FRQs, bond energy calculations, and Hess’s Law problems. Real-world example: Instant cold packs (endothermic) use ammonium nitrate dissolving in water to absorb heat, while hand warmers (exothermic) release heat when iron oxidizes.

Key Terms & Concepts

• Enthalpy (H): Total heat content of a system at constant pressure. ΔH = H_products – H_reactants (change in enthalpy).
• Exothermic reaction: Releases heat to surroundings (ΔH < 0). Example: Combustion of methane (CH₄ + 2O₂ → CO₂ + 2H₂O + heat).
• Endothermic reaction: Absorbs heat from surroundings (ΔH > 0). Example: Photosynthesis (6CO₂ + 6H₂O + light → C₆H₁₂O₆ + 6O₂).
• Standard enthalpy of formation (ΔH°f): Enthalpy change when 1 mole of a compound forms from its elements in standard states (e.g., ΔH°f for H₂O(l) = –285.8 kJ/mol).
• Bond enthalpy (bond energy): Energy required to break 1 mole of bonds in gaseous molecules. ΔH = Σ(bonds broken) – Σ(bonds formed).
• Hess’s Law: The total enthalpy change for a reaction is the same whether it occurs in one step or multiple steps. Used to calculate ΔH for reactions indirectly.
• Calorimetry: Experimental method to measure heat flow using q = mcΔT (q = heat, m = mass, c = specific heat, ΔT = temperature change).
• State function: Property that depends only on initial and final states (e.g., ΔH, ΔG, ΔS), not the path taken.
• Surroundings vs. system: System = the reaction; surroundings = everything else (e.g., water in a calorimeter).
• First Law of Thermodynamics: Energy is conserved (ΔE = q + w, where q = heat, w = work). For reactions at constant pressure, ΔH ≈ ΔE (since work is minimal).

Step-by-Step: Solving Enthalpy Problems

1. Identify the type of problem

• Bond enthalpy? → Use ΔH = Σ(bonds broken) – Σ(bonds formed).
• Formation enthalpies? → Use ΔH°rxn = ΣΔH°f(products) – ΣΔH°f(reactants).
• Calorimetry? → Use q = mcΔT, then relate to moles of reactant.
• Hess’s Law? → Manipulate given reactions to match the target reaction.

2. Write the balanced equation

• Always start with a balanced chemical equation (AP loves testing this!).
• Example: For the combustion of propane (C₃H₈ + 5O₂ → 3CO₂ + 4H₂O), ensure coefficients are correct.

3. Calculate ΔH using the appropriate method

• Bond enthalpy example:
• Break all bonds in reactants (C₃H₈: 2 C–C + 8 C–H; 5O₂: 5 O=O).
• Form all bonds in products (3CO₂: 6 C=O; 4H₂O: 8 O–H).
• ΔH = (energy to break bonds) – (energy to form bonds).
• Formation enthalpy example:
• ΔH°rxn = [3ΔH°f(CO₂) + 4ΔH°f(H₂O)] – [ΔH°f(C₃H₈) + 5ΔH°f(O₂)].
• Note: ΔH°f for elements (e.g., O₂) = 0 kJ/mol.

4. Check units and signs

• ΔH < 0 → Exothermic (heat released).
• ΔH > 0 → Endothermic (heat absorbed).
• Units: kJ/mol (per mole of reaction as written).

5. Apply Hess’s Law (if needed)

• Reverse a reaction? → Flip the sign of ΔH.
• Multiply coefficients? → Multiply ΔH by the same factor.
• Add reactions? → Add their ΔH values.

Common Mistakes

Mistake 1: Forgetting to balance the equation before calculating ΔH.

• Correction: Always balance first! Coefficients affect ΔH (e.g., 2H₂ + O₂ → 2H₂O has twice the ΔH of H₂ + ½O₂ → H₂O).

Mistake 2: Mixing up bond breaking (endothermic) and bond forming (exothermic).

• Correction: Breaking bonds requires energy (ΔH > 0), forming bonds releases energy (ΔH < 0). Think: "Break in, form out."

Mistake 3: Using ΔH°f for elements (e.g., O₂, N₂) in calculations.

• Correction: ΔH°f for elements in their standard states = 0 kJ/mol. Don’t include them in ΣΔH°f(reactants)!

Mistake 4: Ignoring the sign of ΔH in calorimetry problems.

• Correction: If the system loses heat, q_system is negative (exothermic). If the surroundings gain heat, q_surroundings is positive.

Mistake 5: Confusing ΔH with ΔE (internal energy).

• Correction: For reactions at constant pressure, ΔH ≈ ΔE (since work is negligible). At constant volume, ΔE = q_v.

AP Exam Insights

1. FRQs often test Hess’s Law and bond enthalpy.

• Example: Given 3 reactions and their ΔH values, calculate ΔH for a target reaction.
• Trick: Watch for reversed reactions (flip ΔH sign) or scaled coefficients (multiply ΔH).

2. Multiple-choice traps:

• ΔH vs. ΔS vs. ΔG: ΔH is heat, ΔS is entropy (disorder), ΔG is spontaneity. Don’t mix them up!
• Units: kJ vs. kJ/mol. AP often asks for per mole of reaction (e.g., "per mole of CO₂ produced").
• State of matter: ΔH°f for H₂O(l) ≠ H₂O(g). Always check the phase (solid, liquid, gas).

3. Calorimetry problems require careful attention to signs.

• Example: If a reaction in a calorimeter warms the water, the reaction is exothermic (ΔH < 0), but q_water is positive (since it gains heat).

4. Bond enthalpy problems are less precise than ΔH°f.

• Why? Bond enthalpies are averages (e.g., C–H bond energy varies slightly in different molecules). ΔH°f is more accurate.

Quick Check Questions

1. Multiple Choice:

For the reaction: 2H₂(g) + O₂(g) → 2H₂O(l) ΔH = –571.6 kJ What is ΔH for the formation of 1 mole of H₂O(l)? (A) –285.8 kJ (B) –571.6 kJ (C) +285.8 kJ (D) +571.6 kJ

Answer: (A) –285.8 kJ Explanation: The given ΔH is for 2 moles of H₂O. Divide by 2 to get ΔH per mole.

2. Short FRQ:

Given the following data: - ΔH°f for CO₂(g) = –393.5 kJ/mol - ΔH°f for H₂O(l) = –285.8 kJ/mol - ΔH°f for C₃H₈(g) = –103.8 kJ/mol

Calculate ΔH°rxn for the combustion of propane: C₃H₈(g) + 5O₂(g) → 3CO₂(g) + 4H₂O(l)

Answer:
ΔH°rxn = [3(–393.5) + 4(–285.8)] – [–103.8 + 5(0)] = –2219.9 kJ Explanation: Use ΔH°rxn = ΣΔH°f(products) – ΣΔH°f(reactants). ΔH°f for O₂ = 0.

Last-Minute Cram Sheet

1. Exothermic: ΔH < 0 (heat released, feels hot).
2. Endothermic: ΔH > 0 (heat absorbed, feels cold).
3. ΔH = Σ(bonds broken) – Σ(bonds formed) (bond enthalpy).
4. ΔH°rxn = ΣΔH°f(products) – ΣΔH°f(reactants) (formation enthalpy).
5. ΔH°f for elements (e.g., O₂, N₂) = 0 kJ/mol ⚠️.
6. Hess’s Law: Flip reaction → flip ΔH sign; multiply coefficients → multiply ΔH.
7. q = mcΔT (calorimetry: q = heat, m = mass, c = specific heat, ΔT = temp change).
8. State function: ΔH depends only on initial/final states, not the path.
9. Combustion reactions are always exothermic (ΔH < 0).
10. ⚠️ Units matter! kJ vs. kJ/mol (AP often asks for per mole).