---
title: "AP Chem Unit 6 Review: Thermochemistry | Fiveable"
description: "Review AP Chemistry Unit 6 with study guides, practice questions, and key terms on enthalpy, calorimetry, Hess's law, and bond enthalpies."
canonical: "https://fiveable.me/ap-chem/unit-6"
type: "unit"
subject: "AP Chemistry"
unit: "Unit 6 – Thermochemistry"
---

# AP Chem Unit 6 Review: Thermochemistry | Fiveable

## Overview

Unit 6 covers how energy moves between a system and its surroundings during heating, cooling, phase changes, and chemical reactions. You will apply q = mcΔT and q = nΔH for quantitative calculations, read and draw energy diagrams, and calculate reaction enthalpy using bond energies, standard enthalpies of formation, and Hess's law.

## AP CED Alignment

This unit hub is organized around AP Course and Exam Description topics, skills, and exam task types when they are available in the source data.
- 6.1: Endothermic and Exothermic Processes
- 6.2: Energy Diagrams
- 6.3: Heat Transfer and Thermal Equilibrium
- 6.4: Heat Capacity and Calorimetry
- 6.5: Energy of Phase Changes
- 6.6: Introduction to Enthalpy of Reaction
- 6.7: Bond Enthalpies
- 6.8: Enthalpy of Formation
- 6.9: Hess's Law
- 6.3-6.4: Heat Transfer, Thermal Equilibrium, and Calorimetry
- 6.5-6.6: Phase Change Energy and Enthalpy of Reaction
- Practice 5 - Mathematical Routines
- Practice 6 - Argumentation
- Practice 4 - Model Analysis
- FRQ 7 – Short Answer
- FRQ 6 – Short Answer
- FRQ 2 – Long Answer

## Topics

- [6.1: Endothermic and Exothermic Processes](/ap-chem/unit-6/endothermic-exothermic-processes/study-guide/IBke22j0ulvsM9ps3QPJ): Classify chemical and physical changes as endothermic (ΔH > 0, system absorbs energy) or exothermic (ΔH < 0, system releases energy) based on temperature changes in the surroundings and the relative strengths of interactions before and after the process.
- [6.2: Energy Diagrams](/ap-chem/unit-6/energy-diagrams/study-guide/oASenD5gSuLH8VI0Yrxa): Draw and interpret reaction coordinate diagrams showing potential energy versus reaction progress. Identify reactants, products, transition state, activation energy, and the sign of ΔH from the relative heights of reactants and products.
- [6.3: Heat Transfer and Thermal Equilibrium](/ap-chem/unit-6/heat-transfer-thermal-equilibrium/study-guide/WJd1kCvPeSS08sVOaDPt): Explain heat transfer at the particle level: warmer particles have greater average kinetic energy and transfer energy to cooler particles through collisions until thermal equilibrium is reached and temperatures are equal.
- [6.4: Heat Capacity and Calorimetry](/ap-chem/unit-6/heat-capacity-calorimetry/study-guide/jShImkrhZMnPWxlEjdwN): Apply q = mcΔT to calculate heat absorbed or released during temperature changes. Use the first law (q_system + q_surroundings = 0) to solve calorimetry problems involving coffee-cup and bomb calorimeters.
- [6.5: Energy of Phase Changes](/ap-chem/unit-6/energy-phase-changes/study-guide/kAXAzHrD24XL6LdpMFHh): Calculate heat for phase transitions using q = nΔH, where ΔH is the molar enthalpy of fusion or vaporization. Recognize that temperature stays constant during a phase change and that complementary processes (melting/freezing) have equal and opposite ΔH values.
- [6.6: Introduction to Enthalpy of Reaction](/ap-chem/unit-6/intro-enthalpy-reaction/study-guide/MZToO4yTqx0BDiUxmvhN): Connect the sign of ΔH_rxn to the direction of heat flow at constant pressure. Calculate the heat released or absorbed for a given number of moles using q = nΔH_rxn, and scale ΔH when the stoichiometry of the reaction changes.
- [6.7: Bond Enthalpies](/ap-chem/unit-6/bond-enthalpies/study-guide/Q35odLQQyndAgpsA3iwn): Estimate ΔH_rxn by summing average bond energies for all bonds broken in reactants and subtracting the sum for all bonds formed in products. Account for stoichiometric coefficients and recognize that this method gives approximate values for gas-phase reactions.
- [6.8: Enthalpy of Formation](/ap-chem/unit-6/enthalpy-formation/study-guide/glO3L5mcfcUwCd0ODBej): Use tabulated standard enthalpies of formation and the equation ΔH°rxn = ΣΔH°f(products) - ΣΔH°f(reactants) to calculate reaction enthalpy. Remember that ΔH°f = 0 for elements in their standard states and that stoichiometric coefficients must be applied.
- [6.9: Hess's Law](/ap-chem/unit-6/hess-law/study-guide/p9ryCGfaOvpZj0Qye5eT): Apply Hess's law by reversing (flip ΔH sign) and scaling (multiply ΔH by the same factor) thermochemical equations so they add up to a target reaction. Cancel intermediate species algebraically and sum the individual ΔH values to find the overall enthalpy change.

## Hardest Topics And Analytics

Snapshot: practice snapshot
This snapshot uses Fiveable practice activity to show where students tend to miss questions and which review moves are worth prioritizing first.
- **62% average MCQ accuracy** (Across 8.8k multiple-choice practice attempts for this unit.)
- **8.8k MCQ attempts** (Practice activity included in this snapshot.)
- **50% average FRQ score** (Across 24 scored free-response attempts for this unit.)
- **6.7: Bond Enthalpies**: 47% MCQ miss rate across 724 attempts. Review Bond Enthalpies with attention to how the concept appears in AP-style source and evidence questions.
- **6.8: Enthalpy of Formation**: 42% MCQ miss rate across 866 attempts. Review Enthalpy of Formation with attention to how the concept appears in AP-style source and evidence questions.
- **6.5: Energy of Phase Changes**: 40% MCQ miss rate across 1604 attempts. Review Energy of Phase Changes with attention to how the concept appears in AP-style source and evidence questions.
- **6.2: Energy Diagrams**: 36% MCQ miss rate across 1091 attempts. Review Energy Diagrams with attention to how the concept appears in AP-style source and evidence questions.

## Review Notes

### 6.1: Endothermic and Exothermic Processes

Every chemical or physical change either absorbs energy from the surroundings or releases energy to them. The sign of ΔH tells you which direction energy flows: positive ΔH means the system gains energy (endothermic), negative ΔH means the system loses energy (exothermic). Always define the system first before assigning a sign.

- **System vs. surroundings**: The system is the specific matter being studied; the surroundings are everything else. Heat flows between them, and their signs are always opposite.
- **Exothermic process**: Energy leaves the system, so ΔH < 0. The surroundings warm up. Examples include combustion, condensation, and freezing.
- **Endothermic process**: Energy enters the system, so ΔH > 0. The surroundings cool down. Examples include melting, vaporization, and many dissolving processes.
- **Temperature as indicator**: A measurable temperature change in the surroundings confirms that energy transfer occurred. No temperature change during a phase change does not mean no energy transfer.
- **Enthalpy of solution**: Dissolving can be exothermic or endothermic depending on whether the energy released forming solute-solvent interactions is greater or less than the energy required to separate solute and solvent particles.

**Checkpoint:** If a reaction causes the solution in a coffee-cup calorimeter to cool down, is the reaction endothermic or exothermic? Explain using system-surroundings language.

Feature | Exothermic | Endothermic
--- | --- | ---
Sign of ΔH | Negative (ΔH < 0) | Positive (ΔH > 0)
Surroundings temperature | Increases | Decreases
Energy flow | System to surroundings | Surroundings to system
Example process | Combustion, condensation | Melting, many dissolving reactions

### 6.2: Energy Diagrams

An energy diagram (reaction coordinate diagram) plots potential energy on the y-axis against reaction progress on the x-axis. The relative heights of reactants and products show whether the process is endothermic or exothermic, and the peak represents the transition state. For Unit 6, focus on reading ΔH from the diagram rather than activation energy, which is covered more in Unit 5.

- **Exothermic diagram**: Products sit lower than reactants on the y-axis. The energy difference is negative ΔH, and energy is released to the surroundings.
- **Endothermic diagram**: Products sit higher than reactants. The energy difference is positive ΔH, and energy is absorbed from the surroundings.
- **Transition state**: The highest-energy point on the diagram. It represents the activated complex and is not a stable intermediate.
- **Activation energy (Ea)**: The energy difference between the reactants and the transition state. A catalyst lowers Ea without changing the overall ΔH.
- **Phase change diagrams**: Melting and vaporization are shown as uphill steps; freezing and condensation are downhill. The magnitude of the step equals the molar enthalpy of the phase change.

**Checkpoint:** Sketch an energy diagram for an endothermic reaction and label reactants, products, transition state, Ea, and ΔH.

Diagram feature | Exothermic reaction | Endothermic reaction
--- | --- | ---
Product energy vs. reactant energy | Products lower | Products higher
ΔH sign | Negative | Positive
Energy released or absorbed | Released | Absorbed

### 6.3-6.4: Heat Transfer, Thermal Equilibrium, and Calorimetry

Heat transfer occurs at the particle level: warmer particles have greater average kinetic energy and transfer energy to cooler particles through collisions. This continues until thermal equilibrium is reached. Calorimetry quantifies that transfer using q = mcΔT, where q is heat in joules, m is mass in grams, c is specific heat capacity, and ΔT is the temperature change.

- **Thermal equilibrium**: Reached when two bodies in contact have the same average kinetic energy and temperature, so there is no net heat flow between them.
- **q = mcΔT**: The core calorimetry equation. Use it for any process where temperature changes but no phase change occurs. ΔT = T_final - T_initial, so cooling gives negative q.
- **Specific heat capacity (c)**: The energy needed to raise 1 g of a substance by 1°C. Water's value is 4.184 J/g·°C, which is high compared to most metals.
- **First law of thermodynamics**: Energy is conserved. In a calorimetry experiment, q_system + q_surroundings = 0, so heat lost by one substance equals heat gained by the other.
- **Coffee-cup vs. bomb calorimeter**: A coffee-cup calorimeter operates at constant pressure and measures ΔH directly. A bomb calorimeter operates at constant volume and measures ΔE; corrections are needed to get ΔH.

**Checkpoint:** A 50.0 g sample of metal at 95.0°C is placed in 100.0 g of water at 22.0°C. The final temperature is 27.5°C. Calculate the specific heat of the metal using q_metal = -q_water.

Feature | Coffee-cup calorimeter | Bomb calorimeter
--- | --- | ---
Pressure condition | Constant pressure | Constant volume
Quantity measured directly | ΔH | ΔE (internal energy)
Typical use | Dissolution, neutralization | Combustion reactions
Heat equation | q = mcΔT for solution | q = C_cal × ΔT

### 6.5-6.6: Phase Change Energy and Enthalpy of Reaction

During a phase change, temperature stays constant while energy is absorbed or released. The calculation switches from q = mcΔT to q = nΔH, where n is moles and ΔH is the molar enthalpy of the phase transition. The same q = nΔH framework applies to chemical reactions, where ΔH_rxn gives the heat per mole of reaction at constant pressure.

- **q = nΔH for phase changes**: Multiply moles of substance by the molar enthalpy of fusion or vaporization. Melting and boiling are endothermic; freezing and condensation are exothermic with the same magnitude.
- **Temperature plateau**: On a heating or cooling curve, a flat region indicates a phase change. Temperature does not change because added energy breaks intermolecular forces rather than increasing kinetic energy.
- **Enthalpy of vaporization vs. fusion**: ΔH_vap is always larger than ΔH_fus for the same substance because more intermolecular forces must be overcome to convert liquid to gas than solid to liquid.
- **ΔH_rxn sign convention**: Negative ΔH_rxn means the reaction releases heat (exothermic). Positive ΔH_rxn means the reaction absorbs heat (endothermic). Units are kJ/mol of reaction as written.
- **Scaling ΔH_rxn**: If you double the moles of reactant, you double the heat released or absorbed. ΔH is an extensive property when tied to a specific balanced equation.

**Checkpoint:** How much energy is required to vaporize 3.00 mol of water at 100°C if ΔH_vap = 40.7 kJ/mol? Is this process endothermic or exothermic?

Process | Direction | ΔH sign | Equation
--- | --- | --- | ---
Melting (fusion) | Solid to liquid | Positive | q = nΔH_fus
Freezing | Liquid to solid | Negative | q = -nΔH_fus
Vaporization | Liquid to gas | Positive | q = nΔH_vap
Condensation | Gas to liquid | Negative | q = -nΔH_vap

### 6.7: Bond Enthalpies

You can estimate ΔH_rxn by accounting for every bond broken in the reactants and every bond formed in the products. Breaking bonds always requires energy input (endothermic step); forming bonds always releases energy (exothermic step). The net ΔH equals the total energy of bonds broken minus the total energy of bonds formed.

- **ΔH = Σ(bonds broken) - Σ(bonds formed)**: Add up average bond energies for all bonds broken in reactants, then subtract the sum for all bonds formed in products. A positive result means more energy was needed to break bonds than was released forming them.
- **Average bond energy**: A tabulated average value in kJ/mol for a specific bond type (e.g., C-H, O=O, N≡N). Values are averages across many compounds, so this method gives estimates, not exact values.
- **Bond order and strength**: Triple bonds (N≡N, ~945 kJ/mol) are stronger and require more energy to break than double bonds, which require more than single bonds. Shorter bonds are stronger.
- **Counting bonds with coefficients**: Multiply the number of each bond type by the stoichiometric coefficient of that molecule in the balanced equation before summing bond energies.
- **Limitation of bond enthalpies**: Bond energy calculations apply to gas-phase species. Results differ from ΔH°f-based calculations because average bond energies do not account for specific molecular environments.

**Checkpoint:** For the reaction H2(g) + Cl2(g) → 2 HCl(g), use bond energies (H-H: 436, Cl-Cl: 243, H-Cl: 432 kJ/mol) to calculate ΔH and determine whether the reaction is endothermic or exothermic.

### 6.8: Enthalpy of Formation

The standard enthalpy of formation (ΔH°f) is the enthalpy change when exactly 1 mole of a compound forms from its elements in their most stable standard states at 1 bar and 298 K. By definition, ΔH°f = 0 for any element in its standard state. Tabulated ΔH°f values let you calculate ΔH°rxn precisely using the products-minus-reactants formula.

- **ΔH°rxn = ΣΔH°f(products) - ΣΔH°f(reactants)**: Multiply each ΔH°f value by its stoichiometric coefficient, sum the products side, sum the reactants side, and subtract. This gives the standard enthalpy of reaction.
- **Standard state**: The most stable physical form of an element or compound at 1 bar and 298 K. For carbon, graphite is the standard state (ΔH°f = 0), not diamond.
- **ΔH°f = 0 for elements**: Any element in its standard state has ΔH°f = 0 by definition. Examples: O2(g), H2(g), C(graphite), Na(s).
- **State symbols matter**: ΔH°f values differ for different physical states. For example, ΔH°f for H2O(l) and H2O(g) are not the same; always match the state in the balanced equation.
- **Stoichiometric coefficients in summation**: Each ΔH°f value must be multiplied by the coefficient of that species in the balanced equation before summing. Forgetting this is a common calculation error.

**Checkpoint:** Using ΔH°f values, calculate ΔH°rxn for CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(l). Identify which species have ΔH°f = 0.

### 6.9: Hess's Law

Hess's law states that the enthalpy change of an overall reaction equals the sum of the enthalpy changes of any series of steps that add up to that reaction. Because enthalpy is a state function, the path does not matter, only the initial and final states. You manipulate thermochemical equations by reversing them (flip the sign of ΔH) or scaling them (multiply ΔH by the same factor) until they add up to the target reaction.

- **State function property**: Enthalpy depends only on the initial and final states, not the pathway. This is why you can combine any set of steps that give the correct overall equation.
- **Reversing a reaction**: If you reverse a thermochemical equation, the magnitude of ΔH stays the same but the sign changes. An exothermic forward reaction becomes an endothermic reverse reaction.
- **Scaling a reaction**: If you multiply all coefficients by a factor c, multiply ΔH by the same factor c. Halving a reaction halves its ΔH.
- **Canceling intermediates**: Species that appear on both sides of the combined equations cancel out, just like in algebraic addition. The remaining species should match the target equation exactly.
- **Connection to ΔH°f**: The ΔH°f formula is a specific application of Hess's law: you are summing formation reactions for products and subtracting formation reactions for reactants to get the overall reaction enthalpy.

**Checkpoint:** Given: (1) C(s) + O2(g) → CO2(g), ΔH = -393.5 kJ; (2) H2(g) + 1/2 O2(g) → H2O(l), ΔH = -285.8 kJ; (3) CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(l), ΔH = -890.3 kJ. Use Hess's law to find ΔH for C(s) + 2 H2(g) → CH4(g).

Method | What you need | Best used when
--- | --- | ---
Bond enthalpies | Average bond energy table, Lewis structures | Gas-phase reactions, quick estimates
ΔH°f formula | Table of standard formation enthalpies | Standard conditions, precise values needed
Hess's law | Set of related thermochemical equations | Target reaction not directly measurable

## Study Guides

- [6.1 Endothermic and Exothermic Processes](/ap-chem/unit-6/endothermic-exothermic-processes/study-guide/IBke22j0ulvsM9ps3QPJ)
- [6.6 Introduction to Enthalpy of Reaction](/ap-chem/unit-6/intro-enthalpy-reaction/study-guide/MZToO4yTqx0BDiUxmvhN)
- [6.7 Bond Enthalpies](/ap-chem/unit-6/bond-enthalpies/study-guide/Q35odLQQyndAgpsA3iwn)
- [6.3 Heat Transfer and Thermal Equilibrium](/ap-chem/unit-6/heat-transfer-thermal-equilibrium/study-guide/WJd1kCvPeSS08sVOaDPt)
- [6.8 Enthalpy of Formation](/ap-chem/unit-6/enthalpy-formation/study-guide/glO3L5mcfcUwCd0ODBej)
- [6.4 Heat Capacity and Calorimetry](/ap-chem/unit-6/heat-capacity-calorimetry/study-guide/jShImkrhZMnPWxlEjdwN)
- [6.5 Energy of Phase Changes](/ap-chem/unit-6/energy-phase-changes/study-guide/kAXAzHrD24XL6LdpMFHh)
- [6.2 Energy Diagrams ](/ap-chem/unit-6/energy-diagrams/study-guide/oASenD5gSuLH8VI0Yrxa)
- [6.9 Hess’s Law](/ap-chem/unit-6/hess-law/study-guide/p9ryCGfaOvpZj0Qye5eT)

## Practice Preview

### Multiple-choice practice

- **Stimulus-based practice question**: Practice 5 - Mathematical Routines | To calculate the value of $\Delta H_2$ using Hess's Law, which quantities from the diagram must be known?
- **Stimulus-based practice question**: Practice 6 - Argumentation | Which calculation justifies the claim that vaporization requires energy?
- **Stimulus-based practice question**: Practice 6 - Argumentation | Which mathematical justification best supports this claim based on the standard enthalpies of formation shown in the cycle?
- **Stimulus-based practice question**: Practice 4 - Model Analysis | Which of the following best explains the inconsistency with Hess's law?
- **Stimulus-based practice question**: Practice 4 - Model Analysis | Which of the following best explains why the model is consistent with the first law of thermodynamics?
- **Stimulus-based practice question**: Practice 4 - Model Analysis | Which of the following best explains whether this cycle model aligns with chemical theory?

### FRQ practice

- **Combustion thermodynamics and bond enthalpy analysis**: FRQ 7 – Short Answer | Combustion thermodynamics and bond enthalpy analysis
- **Molar enthalpy of combustion of 1-propanol**: FRQ 6 – Short Answer | Molar enthalpy of combustion of 1-propanol
- **Hydrazine combustion enthalpy discrepancies**: FRQ 2 – Long Answer | Hydrazine combustion enthalpy discrepancies

## Key Terms

- **Endothermic Reaction**: A reaction in which the system absorbs heat from the surroundings, giving ΔH > 0. The surroundings decrease in temperature.
- **Exothermic Reaction**: A reaction in which the system releases heat to the surroundings, giving ΔH < 0. The surroundings increase in temperature.
- **Enthalpy Change**: The heat absorbed or released by a system at constant pressure during a chemical or physical process, symbolized ΔH and reported in kJ/mol.
- **Enthalpy of Reaction**: The total heat change for a chemical reaction at constant pressure. Calculated from bond energies, standard enthalpies of formation, or Hess's law.
- **Specific Heat**: The energy required to raise 1 gram of a substance by 1°C. Water's specific heat is 4.184 J/g·°C, used in q = mcΔT calculations.
- **Thermal Equilibrium**: The state reached when two objects in thermal contact have the same temperature and there is no net heat flow between them.
- **calorimeter**: An insulated device used to measure heat transfer during a chemical or physical process. Coffee-cup calorimeters operate at constant pressure; bomb calorimeters operate at constant volume.
- **First Law Of Thermodynamics**: Energy is conserved in all chemical and physical processes. In calorimetry, q_system + q_surroundings = 0.
- **Enthalpy of Fusion**: The energy required to melt 1 mole of a solid at its melting point at constant pressure. The reverse process (freezing) releases the same amount of energy.
- **enthalpy of vaporization**: The energy required to convert 1 mole of liquid to gas at constant temperature and pressure. Always larger than the enthalpy of fusion for the same substance.
- **Energy Diagram**: A graph of potential energy versus reaction progress showing reactant and product energy levels, the transition state, activation energy, and the sign of ΔH.
- **Potential Energy**: Stored energy in chemical bonds and intermolecular forces. Bond breaking increases potential energy; bond formation decreases it.
- **Surroundings**: Everything outside the system being studied. Heat flows from system to surroundings in exothermic processes and from surroundings to system in endothermic processes.
- **Stoichiometric Coefficients**: The numbers in a balanced equation that indicate molar ratios. In thermochemistry, each ΔH°f or bond energy value must be multiplied by the corresponding coefficient.
- **Average Kinetic Energy**: The average energy of particle motion in a substance, directly proportional to temperature. At thermal equilibrium, two bodies have the same average kinetic energy.

## Common Mistakes

- **Mixing up the sign of ΔH for the system vs. surroundings**: Students often assign the sign based on what happens to the surroundings rather than the system. If the surroundings warm up, the system lost energy, so ΔH for the system is negative (exothermic). Always define the system first.
- **Forgetting stoichiometric coefficients in bond enthalpy and ΔH°f calculations**: Each bond energy or ΔH°f value must be multiplied by the coefficient of that species in the balanced equation. Skipping this step is the most common arithmetic error in Unit 6 calculations.
- **Using q = mcΔT during a phase change**: Temperature does not change during a phase transition, so q = mcΔT gives zero and is the wrong equation. Use q = nΔH for any process occurring at a temperature plateau on a heating or cooling curve.
- **Not flipping the sign when reversing a reaction in Hess's law**: When you reverse a thermochemical equation to build a target reaction, the sign of ΔH must change. Forgetting this step leads to an answer with the wrong sign or wrong magnitude.
- **Treating ΔH°f = 0 only for monatomic elements**: ΔH°f = 0 applies to any element in its most stable standard state, including diatomic molecules like O2(g), H2(g), N2(g), and Cl2(g), as well as graphite for carbon. Using diamond or O3 as the standard state for carbon or oxygen is incorrect.

## Exam Connections

- **Multi-step calculation tasks**: AP Chemistry free-response questions in this unit frequently require chained calculations: for example, using q = mcΔT to find heat from calorimetry data, then converting to kJ/mol of reactant. Showing each step with correct units and sign conventions is essential for full credit.
- **Justification using particle-level reasoning**: Multiple-choice and free-response questions often ask you to explain why a process is endothermic or exothermic in terms of bond breaking and forming, or to connect a temperature change in the surroundings to the direction of heat flow. Answers that only state the sign of ΔH without a mechanism typically earn partial credit at best.
- **Selecting and applying the correct enthalpy calculation method**: The exam tests all three methods for finding ΔH_rxn: bond enthalpies, standard enthalpies of formation, and Hess's law. Questions may provide data suited to one method and require you to recognize which approach applies, manipulate the given equations or values correctly, and interpret the result in context.

## Final Review Checklist

- **Classify processes and assign ΔH signs**: For any chemical or physical change, identify the system and surroundings, determine whether the process is endothermic or exothermic from temperature observations, and assign the correct sign to ΔH.
- **Draw and read energy diagrams**: Sketch a reaction coordinate diagram for both endothermic and exothermic processes. Label reactants, products, transition state, Ea, and ΔH. Read ΔH as the energy difference between products and reactants.
- **Apply q = mcΔT and q = nΔH correctly**: Use q = mcΔT for temperature changes and q = nΔH for phase changes. In calorimetry problems, set q_system = -q_surroundings and watch the sign of ΔT.
- **Calculate ΔH using bond enthalpies**: Sum bond energies for all bonds broken in reactants, sum bond energies for all bonds formed in products, and subtract. Multiply by stoichiometric coefficients and apply the correct sign convention.
- **Use the ΔH°f formula with correct coefficients**: Apply ΔH°rxn = ΣΔH°f(products) - ΣΔH°f(reactants), multiplying each ΔH°f by its stoichiometric coefficient. Confirm that elements in standard states have ΔH°f = 0 and match physical state symbols.
- **Manipulate thermochemical equations with Hess's law**: Reverse equations (flip ΔH sign) and scale equations (multiply ΔH by the same factor) to build a target reaction. Cancel intermediate species and sum all ΔH values to get the overall enthalpy change.
- **Connect Unit 6 to Unit 9**: Recognize that ΔH from Unit 6 is one component of Gibbs free energy (ΔG = ΔH - TΔS) in Unit 9. A strongly negative ΔH favors spontaneity but does not guarantee it without considering entropy.

## Study Plan

- **Step 1: Build the conceptual foundation (Topics 6.1-6.3)**: Read the topic guides for 6.1, 6.2, and 6.3. Practice classifying processes as endothermic or exothermic using system-surroundings language. Sketch energy diagrams for both types and label all components. Review the particle-level explanation of heat transfer and thermal equilibrium.
- **Step 2: Practice calorimetry calculations (Topics 6.3-6.4)**: Work through q = mcΔT problems with both coffee-cup and bomb calorimeter setups. Practice applying q_system + q_surroundings = 0 to find unknown specific heats or temperature changes. Check sign conventions carefully for each calculation.
- **Step 3: Understand phase change and reaction enthalpy (Topics 6.5-6.6)**: Practice switching between q = mcΔT and q = nΔH on multi-step heating curve problems. Then apply q = nΔH_rxn to chemical reactions, scaling ΔH when the moles of reactant change. Use the topic guides for 6.5 and 6.6 to review worked examples.
- **Step 4: Work all three ΔH calculation methods (Topics 6.7-6.8)**: Complete bond enthalpy problems using the ΔH = Σ(bonds broken) - Σ(bonds formed) formula, then practice the ΔH°f formula with the products-minus-reactants equation. Compare results from both methods for the same reaction to understand why they differ.
- **Step 5: Apply Hess's law and review the full unit (Topic 6.9)**: Practice reversing and scaling thermochemical equations to build target reactions. Then use the AP score calculator to estimate your estimated score range and identify which calculation type needs the most additional practice before the exam.

## More Ways To Review

- [Topic study guides](/ap-chem/unit-6#topics)
- [FRQ practice](/ap-chem/frq-practice)
- [Cram archive videos](/cram-archives?subject=ap-chemistry&unit=unit-6)
- [Cheatsheets](/ap-chem/cheatsheets/unit-6)
- [Key terms](/ap-chem/key-terms)

## FAQs

### What topics are covered in AP Chem Unit 6?

AP Chem Unit 6 covers 9 topics in thermochemistry: Endothermic and Exothermic Processes, Energy Diagrams, Heat Transfer and Thermal Equilibrium, Heat Capacity and Calorimetry, Energy of Phase Changes, Introduction to Enthalpy of Reaction, Bond Enthalpies, Enthalpy of Formation, and Hess's Law. Together they build a complete picture of how energy moves in chemical and physical processes. See [AP Chem Unit 6](/ap-chem/unit-6) for matched practice on each topic.

### How much of the AP Chem exam is Unit 6?

AP Chem Unit 6 makes up 7-9% of the AP exam. That slice covers thermochemistry concepts including heat capacity, calorimetry, enthalpy of reaction, bond enthalpies, enthalpy of formation, and Hess's Law. It's a focused unit, so strong performance here is very achievable with targeted practice.

### What's on the AP Chem Unit 6 progress check (MCQ and FRQ)?

The AP Chem Unit 6 progress check includes both MCQ and FRQ parts drawn from all 9 thermochemistry topics. MCQ questions test concepts like endothermic vs. exothermic processes, heat capacity, calorimetry calculations, and energy diagrams. FRQ questions typically ask you to calculate enthalpy changes using Hess's Law, bond enthalpies, or enthalpy of formation data. Practicing these topics before the progress check is the best prep. Find matched questions at [AP Chem Unit 6](/ap-chem/unit-6).

### How do I practice AP Chem Unit 6 FRQs?

AP Chem Unit 6 FRQs most often ask you to calculate enthalpy changes using Hess's Law, enthalpy of formation tables, or bond enthalpies, and to interpret calorimetry data using heat capacity equations. To practice, work through multi-step calculation problems where you show each step clearly, since College Board awards points for process, not just the final answer. Start with topic-level practice at [AP Chem Unit 6](/ap-chem/unit-6), focusing on Topics 6.6 through 6.9 where FRQ prompts are most common.

### Where can I find AP Chem Unit 6 practice questions?

For AP Chem Unit 6 practice questions, including multiple-choice and practice test style problems, head to [AP Chem Unit 6](/ap-chem/unit-6). You'll find MCQ sets covering heat capacity, calorimetry, energy diagrams, and enthalpy calculations, plus FRQ practice for Hess's Law and enthalpy of formation. Working through both question types gives you the best coverage of the 7-9% exam weight this unit carries.

### How should I study AP Chem Unit 6?

Start AP Chem Unit 6 by making sure you can identify endothermic and exothermic processes and read energy diagrams before moving to calculations. Then build your heat capacity and calorimetry skills, since those equations show up in both MCQ and FRQ. Once calculations feel solid, work through enthalpy of reaction, bond enthalpies, enthalpy of formation, and Hess's Law in order, because each topic builds on the last. A few concrete steps that help:
- Write out Hess's Law problems by hand until flipping and scaling equations feels automatic.
- Practice calorimetry problems with real data sets, not just plug-and-chug examples.
- Review energy diagrams for both phase changes and reactions side by side. Find topic-by-topic practice at [AP Chem Unit 6](/ap-chem/unit-6).

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