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AP Chem Unit 6 Review: Thermochemistry

Q: 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 for matched practice on each topic.

Q: 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.

Q: 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 .

Q: 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 , focusing on Topics 6.6 through 6.9 where FRQ prompts are most common.

Q: 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 . 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.

Q: 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 .

Review AP Chem Unit 6 to build a working understanding of energy changes in chemical and physical processes, from heat transfer and calorimetry to bond enthalpies, enthalpy of formation, and Hess's law. This unit carries 7-9% of the AP exam and connects directly to thermodynamics in Unit 9.

Fiveable's AP Chemistry Unit 6 review covers 9 topics, with guided practice questions for the unit.

Use the topic guides, key terms, and available practice questions to work through every concept from endothermic and exothermic processes through Hess's law calculations.

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AP exam review verified for 2027

What is AP Chem unit 6?

Thermochemistry is the study of heat flow in chemical and physical processes. Every reaction either releases energy to the surroundings or absorbs energy from them, and the first law of thermodynamics guarantees that energy is conserved throughout. Unit 6 gives you the vocabulary, diagrams, and calculation methods to quantify those energy changes.

Unit 6 asks you to classify processes as endothermic or exothermic, calculate heat using q = mcΔT and q = nΔH, interpret energy diagrams, and find reaction enthalpy three ways: from bond energies, from standard enthalpies of formation, and by applying Hess's law.

Energy and temperature

A temperature increase in the surroundings signals an exothermic process (ΔH < 0); a temperature decrease signals an endothermic process (ΔH > 0). Topics 6.1-6.3 build this conceptual foundation using system-versus-surroundings language and molecular collision reasoning.

Quantitative heat calculations

Topics 6.4 and 6.5 introduce the two core equations: q = mcΔT for temperature changes and q = nΔH for phase changes. Calorimetry experiments use these equations together, and the first law requires that q_system + q_surroundings = 0.

Three methods for reaction enthalpy

Topics 6.6-6.9 give you three calculation routes: using q = nΔH_rxn directly from calorimetry data, estimating ΔH from average bond energies (bonds broken minus bonds formed), applying the formation equation ΔH°rxn = ΣΔH°f(products) - ΣΔH°f(reactants), and combining thermochemical equations with Hess's law.

Energy is conserved, and its direction of flow determines reaction character

The first law of thermodynamics runs through every topic in Unit 6. Whether you are dissolving a salt, melting ice, or combusting methane, the total energy of the universe does not change. Thermochemistry gives you the tools to track where that energy goes and how much of it moves, which is the foundation for predicting reaction favorability in Unit 9.

AP Chem unit 6 topics

6.1

Endothermic and Exothermic Processes

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.

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6.2

Energy Diagrams

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.

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6.3

Heat Transfer and Thermal Equilibrium

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.

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6.4

Heat Capacity and Calorimetry

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.

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6.5

Energy of Phase Changes

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.

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6.6

Introduction to Enthalpy of Reaction

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.

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6.7

Bond Enthalpies

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.

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6.8

Enthalpy of Formation

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.

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6.9

Hess's Law

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.

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practice snapshot

Hardest AP Chemistry unit 6 topics

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.

8.8kMCQ attempts

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50%average FRQ score

Across 24 scored free-response attempts for this unit.

Hardest topics in unit 6

MCQ miss rate

6.7

Bond Enthalpies

Review Bond Enthalpies with attention to how the concept appears in AP-style source and evidence questions.

47%724 tries

6.8

Enthalpy of Formation

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42%866 tries

6.5

Energy of Phase Changes

Review Energy of Phase Changes with attention to how the concept appears in AP-style source and evidence questions.

40%1,604 tries

6.2

Energy Diagrams

Review Energy Diagrams with attention to how the concept appears in AP-style source and evidence questions.

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Unit 6 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.

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.

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

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.

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

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.

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.

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.

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.

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

Practice AP Chem unit 6 questions

Try stimulus-based AP practice questions and written prompts after you review the notes.

Example stimulus-based MCQs

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process_diagram

Stimulus-based practice question

A student is analyzing the thermochemistry of nitrogen oxides. The diagram represents an enthalpy cycle connecting nitrogen and oxygen gas to nitrogen dioxide and dinitrogen tetroxide, with arrows indicating specific reaction pathways.

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Question

To calculate the value of $\Delta H_2$ using Hess's Law, which quantities from the diagram must be known?

The values of $\Delta H_3$ and $\Delta H_1$ only, because $\Delta H_2$ equals $\Delta H_3$ plus $\Delta H_1$ .

The value of $\Delta H_1$ only, because $\Delta H_2$ is the exact reverse of the $\Delta H_1$ step.

The values of $\Delta H_3$ and $\Delta H_1$ only, because $\Delta H_2$ equals $\Delta H_3$ minus $\Delta H_1$ .

The value of $\Delta H_3$ only, because $\Delta H_2$ is exactly equal to the $\Delta H_3$ pathway.

bar_chart

Stimulus-based practice question

For the phase change $H_2O(l) \rightarrow H_2O(g)$ , standard enthalpies of formation for $H_2O(l)$ and $H_2O(g)$ are provided.

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Question

Which calculation justifies the claim that vaporization requires energy?

$\Delta H^\circ = -242 - (-286) = -44\ \text{kJ/mol}$ because vaporization is exothermic.

$\Delta H^\circ = -242 + (-286) = -528\ \text{kJ/mol}$ because vaporization is endothermic.

$\Delta H^\circ = -242 - (-286) = +44\ \text{kJ/mol}$ because vaporization releases energy.

$\Delta H^\circ = -242 - (-286) = +44\ \text{kJ/mol}$ because vaporization is endothermic.

Example FRQs

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SAQ

Combustion thermodynamics and bond enthalpy analysis

7. A scientist investigates the thermodynamics of the combustion of liquid ethanol, $C_2H_5OH(l)$ . The reaction proceeds according to the balanced equation below.

$C_2H_5OH(l) + 3 O_2(g) \rightarrow 2 CO_2(g) + 3 H_2O(l)$

The scientist generates an energy diagram for the reaction, which is shown in Figure 1.

Figure 1. Potential energy diagram for the combustion of liquid ethanol, showing the reactant energy level at 0 kJ/mol, the transition-state maximum at 150 kJ/mol, and the product energy level at −1370 kJ/mol.

Calculate the activation energy, $E_a$ , in kJ/mol, for the forward reaction shown in Figure 1.

Table 1. Standard Enthalpies of Formation

Substance	Standard Enthalpy of Formation, $\Delta H^\circ_f$ (kJ/mol)
$C_2H_5OH(l)$	-277.7
$CO_2(g)$	-393.5
$H_2O(l)$	-285.8

The scientist finds standard enthalpy of formation data for the substances involved in the reaction, which are listed in Table 1.

Calculate the standard enthalpy change, $\Delta H^\circ_{rxn}$ , for the combustion of ethanol using the data in Table 1.

ii.

Using your answer to part B(i), calculate the amount of heat released, in kJ, when a 5.00 g sample of $C_2H_5OH(l)$ (molar mass 46.07 g/mol) is burned completely.

A student claims that the bonds in the product molecules are weaker than the bonds in the reactant molecules because energy was released during the reaction. Do you agree or disagree? Justify your answer based on the principles of bond enthalpy.

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SAQ

Molar enthalpy of combustion of 1-propanol

6. A student conducts an experiment to determine the molar enthalpy of combustion of 1-propanol, $C_3H_7OH(l)$ . The student uses the calorimetry setup shown in Figure 1 and records the data shown in Table 1 during the experiment.

Table 1. Experimental data for the combustion of 1-propanol

Measurement	Value
Mass of water in beaker	150.0 g
Initial temperature of water	22.0 °C
Final temperature of water	44.5 °C
Initial mass of burner and fuel	185.45 g
Final mass of burner and fuel	184.82 g

Figure 1. Calorimetry experimental setup for determining the molar enthalpy of combustion of 1-propanol

Write the balanced chemical equation for the complete combustion of 1-propanol, $C_3H_7OH(l)$ , to produce $CO_2(g)$ and $H_2O(l)$ .

Calculate the amount of heat, $q$ , in Joules, absorbed by the water in the beaker. Assume the specific heat capacity of water is $4.18 \text{ J}/(g·^\circ\text{C})$ .

Calculate the experimental molar enthalpy of combustion, $\Delta H_{comb}$ , of 1-propanol in kJ/mol. The molar mass of 1-propanol is 60.10 g/mol.

Table 2. Thermochemical data for selected fuels

Fuel Name	Formula	Molar Mass (g/mol)	Standard Enthalpy of Combustion ( $\Delta H^{\circ}_{comb}$ ) (kJ/mol)
Methanol	$CH_3OH$	32.04	-726
Ethanol	$C_2H_5OH$	46.07	-1368
Octane	$C_8H_{18}$	114.23	-5470

Calculate the mass of each fuel in Table 2 required to produce 500. kJ of heat, and identify the fuel that requires the smallest mass. The student considers using other fuels for a different application where minimizing the mass of fuel carried is critical. Data for three potential fuels are provided in Table 2.

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FRQ

Hydrazine combustion enthalpy discrepancies

2. Answer the following questions about hydrazine, N₂H₄.

Hydrazine, N₂H₄(l), is a liquid fuel used in rocket propulsion. It reacts with oxygen gas, O₂(g), to produce nitrogen gas and water. A student investigates the enthalpy of combustion of hydrazine using the experimental setup shown in Figure 1.

N₂H₄(l) + O₂(g) → N₂(g) + 2H₂O(l)

In the experiment, the student heats 150.0 g of water in the metal can. The initial temperature of the water is 22.0°C, and the final temperature is 32.8°C. The mass of the N₂H₄(l) burned is 0.350 g. Assume the specific heat capacity of the water is 4.18 J/(g·°C) and that the heat capacity of the metal can is negligible.

Calculate the amount of heat, q, in Joules, absorbed by the water.

ii.

Calculate the experimental molar enthalpy of combustion, ΔH_comb, of hydrazine in kJ/mol_rxn.

The student considers the gas-phase reaction represented by the equation below. Lewis electron-dot diagrams for the species involved are shown in Figure 2, and average bond enthalpies are provided in Table 1.

N₂H₄(g) + O₂(g) → N₂(g) + 2H₂O(g)

Figure 2. Lewis diagrams for reactants and products in hydrazine combustion

Table 1. Average Bond Enthalpies

Bond	Bond Enthalpy (kJ/mol)
N–H	391
N–N	163
N≡N	941
O=O	495
O–H	463

Using the Lewis diagrams in Figure 2 and the bond enthalpies in Table 1, calculate the approximate enthalpy change, ΔH, for the gas-phase reaction N₂H₄(g) + O₂(g) → N₂(g) + 2H₂O(g).

The value calculated in part B (-571 kJ/mol) is less negative than the accepted standard enthalpy of combustion for liquid hydrazine (-622 kJ/mol). Explain the thermodynamic reason for this difference in terms of the states of matter of the reactants or products.

The standard enthalpy of formation, ΔH°f, for H₂O(l) is -286 kJ/mol. The standard enthalpy of formation for N₂(g) is 0 kJ/mol.

Using the accepted enthalpy of combustion for the reaction N₂H₄(l) + O₂(g) → N₂(g) + 2H₂O(l) (ΔH°_comb = -622 kJ/mol) and the standard enthalpies of formation given above, calculate the standard enthalpy of formation, ΔH°f, of N₂H₄(l).

An energy diagram for the combustion of hydrazine is shown in Figure 3.

Figure 3. Potential energy diagram for the combustion reaction of hydrazine

Based on the energy diagram in Figure 3, is the combustion of hydrazine an exothermic or endothermic process? Justify your answer.

Figure 1. Calorimetry setup used to measure heat released from combustion of hydrazine

In a second trial, the student uses a glass beaker instead of the metal can shown in Figure 1. The student observes that the magnitude of the calculated ΔH_comb is significantly lower than in the first trial. Explain this observation.

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AP Chem unit 6 practice test

A timed, AP-style test built only from unit 6 questions — multiple choice plus free response, scored with AI feedback. Use the standard format or build your own.

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Key terms

Term	Definition
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 unit 6 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.

How this unit shows up on the AP exam

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 unit 6 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.

How to study unit 6

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

Open the individual guides for Unit 6 when you want a closer review of one topic.

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Practice questions

Use AP-style practice after you review the notes so you can check what you understand.

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FRQ practice

Practice free-response reasoning and compare your answer with scoring guidance.

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Cram archive videos

Watch past review streams filtered to Unit 6 when you want a video walkthrough.

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Official unit cheatsheet

Open the Fiveable one-page unit review, then explore visual cheatsheets for a quick refresher.

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Score calculator

Estimate your broader AP score goal after you review the course and exam format.

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Unit 6 printables

Print the unit cheatsheets ($3) or full-length exam ($12) for Unit 6. Yours to keep, no subscription needed.

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Frequently Asked Questions

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 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.

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, 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. 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.

Ready to review Unit 6?Start with the notes, check the topic cards, and use the practice or resource links when they are available for this course.

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