Energy is at the heart of chemistry. Thermochemistry is the study of energy changes associated with chemical reactions and physical processes. Whether a reaction releases heat or absorbs it, and how much, determines whether it's favorable and practical. Understanding thermochemistry lets you predict what happens in a system and calculate the energy involved.
🎯 What You Need to Know for the Exam
Unit 6 makes up about 7-9% of the AP Chemistry exam. Focus your energy on these priorities:
- Distinguishing between endothermic and exothermic reactions and recognizing observable evidence of each
- Using calorimetry and the heat equation (q = mcΔT) to calculate energy changes from experimental data
- Understanding molar enthalpy of formation and using it to calculate reaction enthalpy
- Calculating enthalpy changes using bond energies and Hess's Law
- Understanding heat capacity, specific heat, and molar heat capacity and their roles in energy calculations
- Connecting energy diagrams to endothermic and exothermic processes
What's in this review:
Topic 6.1: Endothermic and Exothermic Processes
Every chemical reaction and physical process involves energy. An exothermic process releases energy to the surroundings as heat. An endothermic process absorbs energy from the surroundings.
When a reaction is exothermic, the products have lower energy than the reactants. The energy difference is released, typically as heat, warming the surroundings. You observe exothermic reactions as temperature increases. When reactants burn, they're exothermic. Hand warmers are exothermic chemical reactions.
An endothermic process does the opposite. Products have higher energy than reactants. The reaction absorbs energy from the surroundings, cooling them. Melting ice is endothermic. Dissolving some salts in water is endothermic and the solution feels cold.
The first law of thermodynamics states that energy is conserved. In an exothermic reaction, the energy lost by the chemical system is gained by the surroundings. In an endothermic reaction, the system gains energy and the surroundings lose it.
Key concepts to know:
- Exothermic: Releases energy to surroundings. Products lower energy than reactants. Temperature of surroundings increases.
- Endothermic: Absorbs energy from surroundings. Products higher energy than reactants. Temperature of surroundings decreases.
- Energy conservation: Total energy is conserved. System energy change equals negative of surroundings energy change.
- Observable evidence: Combustion, neutralization, and precipitation are usually exothermic. Phase changes like melting and evaporation are endothermic.
⚠ Watch out for:
Exothermic doesn't mean "fast" and endothermic doesn't mean "slow." Energy type and reaction rate are independent. A reaction can be fast and endothermic, or slow and exothermic. Don't confuse them.
Topic 6.2: Energy Diagrams
An energy diagram shows the energy of a system at different stages of a process. For a chemical reaction, it shows the energy of reactants on the left, the energy of products on the right, and the change in between.
If products are lower in energy than reactants, the diagram shows a downward slope, indicating an exothermic process. If products are higher, the diagram slopes upward, indicating an endothermic process. The height difference is the enthalpy change (ΔH) of the reaction.
Energy diagrams make it easy to compare exothermic and endothermic processes at a glance. They're often used to show whether a reaction releases or absorbs energy, and by how much.
Key concepts to know:
- Energy diagram: Shows energy of reactants, products, and the change between them.
- Exothermic: Slopes downward from left to right. Products lower energy. ΔH is negative.
- Endothermic: Slopes upward from left to right. Products higher energy. ΔH is positive.
⚠ Watch out for:
Simple energy diagrams show only initial and final states, not the reaction pathway. Don't confuse these with reaction energy profiles (which show activation energy and transition states).
Topic 6.3: Heat Transfer and Thermal Equilibrium
Heat is the transfer of energy due to temperature difference. When two objects at different temperatures are in contact, particles collide and energy transfers from the warmer object (with higher average kinetic energy) to the cooler object (with lower average kinetic energy).
This process continues until thermal equilibrium is reached. At thermal equilibrium, both objects have the same temperature and the same average kinetic energy per particle. Macroscopically, no net heat flows between them anymore.
The rate and amount of heat transfer depend on the temperature difference. A larger difference drives faster energy transfer. Eventually, equilibrium is established.
Key concepts to know:
- Heat transfer: Energy flow due to temperature difference. Driven by differences in particle kinetic energy.
- Thermal equilibrium: Both bodies reach the same temperature. Average kinetic energy per particle is equal. No net heat flow.
- Particle collisions: Direct mechanism of heat transfer. Faster particles transfer energy to slower particles.
⚠ Watch out for:
Heat is energy in transit due to temperature difference. Temperature is the average kinetic energy of particles. These are related but not the same thing. An object doesn't "contain heat"; it contains thermal energy.
Topic 6.4: Heat Capacity and Calorimetry
The heat equation (q = mcΔT) is one of the most important formulas in thermochemistry. It calculates heat absorbed or released as a substance is heated or cooled. Here, q is heat, m is mass, c is specific heat capacity, and ΔT is the temperature change.
Specific heat capacity is the amount of heat required to raise the temperature of 1 gram of a substance by 1 degree Celsius. Water has a high specific heat (4.18 J/g°C), meaning it takes a lot of energy to change its temperature. This is why oceans moderate climate. Molar heat capacity is the energy per mole per degree.
Calorimetry is the experimental technique for measuring heat using a calorimeter. When a reaction occurs in a calorimeter, the heat released or absorbed changes the temperature of the surrounding water. By measuring the temperature change and knowing the mass and specific heat of water, you calculate the energy involved in the reaction.
The first law of thermodynamics appears here: energy is conserved. Heat released by the reaction equals heat absorbed by the water. Energy lost by one equals energy gained by the other.
Key concepts to know:
- Heat equation: q = mcΔT. Heat = mass × specific heat × temperature change.
- Specific heat capacity: Energy per unit mass per degree temperature change.
- Calorimetry: Experimental measurement of heat using temperature change in water.
- Energy conservation: Heat released by reaction = heat absorbed by surroundings.
⚠ Watch out for:
Watch your signs. If a reaction heats water, ΔT is positive and q is positive (heat absorbed). If a reaction cools water, ΔT is negative and q is negative (heat released). Also, heat released by reaction and heat absorbed by water have opposite signs.
Topic 6.5: Energy of Phase Changes
Phase changes (melting, freezing, vaporization, condensation) involve energy but no temperature change. During a phase change, the temperature of a pure substance stays constant. All the energy goes into rearranging molecules, not increasing their motion.
Molar enthalpy of fusion (ΔHfus) is the energy required to melt 1 mole of solid to liquid at constant pressure. Molar enthalpy of vaporization (ΔHvap) is the energy required to vaporize 1 mole of liquid to gas. Vaporization requires more energy than fusion because molecules in the gas phase are much more separated.
Complementary phase changes have opposite energy changes. The energy to melt solid into liquid equals the negative of the energy to freeze liquid back to solid. Similarly, energy to vaporize equals the negative of energy to condense.
To calculate energy in a phase change: q = n × ΔH, where n is moles and ΔH is molar enthalpy of the phase change.
Key concepts to know:
- Molar enthalpy of fusion: Energy to melt 1 mole of solid. Constant temperature during phase change.
- Molar enthalpy of vaporization: Energy to vaporize 1 mole of liquid. Also constant temperature.
- Complementary relationships: Fusion and freezing have equal magnitude but opposite signs. Same for vaporization and condensation.
- Heat calculation: q = n × ΔH for phase changes.
⚠ Watch out for:
During a phase change, temperature doesn't change, so don't use q = mcΔT. Use q = n × ΔH instead. Also, vaporization always requires more energy than fusion for the same substance.
Topic 6.6: Introduction to Enthalpy of Reaction
The enthalpy change (ΔH) of a reaction is the heat absorbed or released at constant pressure. This is what you measure in a calorimeter at atmospheric pressure. Exothermic reactions have negative ΔH (heat released). Endothermic reactions have positive ΔH (heat absorbed).
The sign and magnitude of ΔH tell you whether a reaction is practical. A large negative ΔH means the reaction releases a lot of energy. The exact value depends on how much reactant reacts, which is why ΔH is usually reported per mole of reaction.
When you write a thermochemical equation, include ΔH. For example: 2 H2 + O2 → 2 H2O, ΔH = -571.6 kJ. This tells you that burning 2 moles of hydrogen releases 571.6 kJ.
Key concepts to know:
- Enthalpy change (ΔH): Heat absorbed or released at constant pressure. Unit: kJ/mol or similar.
- Exothermic: Negative ΔH. Heat released to surroundings.
- Endothermic: Positive ΔH. Heat absorbed from surroundings.
- Thermochemical equations: Include ΔH to show energy change for the reaction as written.
⚠ Watch out for:
ΔH depends on the quantities in the equation. If you reverse the equation, ΔH changes sign. If you double the coefficients, ΔH doubles. Always pay attention to what equation the ΔH value corresponds to.
Topic 6.7: Bond Enthalpies
Chemical reactions involve breaking bonds in reactants and forming bonds in products. Breaking bonds requires energy (endothermic). Forming bonds releases energy (exothermic). The net energy change of the reaction is the difference.
Bond enthalpy is the average energy required to break a bond. Using bond enthalpies, you can estimate reaction enthalpies without measuring them experimentally. The calculation is: ΔH = (energy to break bonds) - (energy to form bonds).
This method is an approximation because bond enthalpies are averages. The exact energy depends on the molecular environment. Still, bond enthalpy calculations give good estimates.
Key concepts to know:
- Bond enthalpy: Average energy to break a bond. Always positive (requires energy).
- Breaking bonds: Endothermic. Contributes positive energy to ΔH.
- Forming bonds: Exothermic. Contributes negative energy to ΔH.
- Calculation: ΔH = Σ(bond enthalpies of bonds broken) - Σ(bond enthalpies of bonds formed).
⚠ Watch out for:
Don't forget to count all bonds. If a reaction breaks 2 C-H bonds, include both. Also, bond enthalpies are averages, so calculations give estimates, not exact values.
Topic 6.8: Enthalpy of Formation
The standard enthalpy of formation (ΔH°f) is the enthalpy change when 1 mole of a compound forms from its elements in their standard states. For example, ΔH°f for water is the enthalpy of: H2(g) + 1/2 O2(g) → H2O(l).
Elements in their standard states have ΔH°f = 0 by definition. You can't calculate their formation enthalpy because there's no reaction forming them from simpler substances.
To find the enthalpy change of any reaction, use: ΔH°reaction = Σ(ΔH°f of products) - Σ(ΔH°f of reactants). Tables of standard enthalpies of formation let you calculate reaction enthalpies without doing experiments.
This method is more accurate than bond enthalpies because it's based on actual measured values.
Key concepts to know:
- Enthalpy of formation: ΔH°f for forming 1 mole of compound from elements in standard states.
- Standard states: Typically most stable form at 25°C and 1 atm. ΔH°f for elements = 0.
- Reaction calculation: ΔH° = Σ(ΔH°f products) - Σ(ΔH°f reactants).
- Tables: Use standard enthalpy tables for accurate calculations.
⚠ Watch out for:
Pay attention to stoichiometry. If an equation has 2 moles of a compound, multiply its ΔH°f by 2. Also, ΔH°f depends on the physical state (s, l, g) of the compound. H2O(l) and H2O(g) have different values.
Topic 6.9: Hess's Law
Hess's Law states that the enthalpy change of a reaction is the same whether it happens in one step or many steps. Because enthalpy depends only on the initial and final states, and not on the path taken, you can combine simpler reactions to find the enthalpy of a more complex one.
Because of this property, you can add chemical equations together. When you add equations, you also add their ΔH values. If you reverse an equation, the sign of ΔH flips. If you multiply an equation by a factor, ΔH multiplies by the same factor.
This allows you to calculate the enthalpy of any reaction by combining simpler reactions with known ΔH values. For example, if you know the reactions for forming CO2 and H2O, you can calculate the enthalpy of combustion of a hydrocarbon.
Key concepts to know:
- Hess's Law: ΔH for a reaction is the sum of ΔH values for individual steps.
- Path independence: ΔH depends only on initial and final states, not the path taken between them.
- Manipulation rules: Reverse an equation, flip the sign of ΔH. Multiply an equation, multiply ΔH by the same factor.
- Adding equations: Sum the equations and sum their ΔH values to find the overall reaction and its enthalpy.
⚠ Watch out for:
When you reverse an equation, reverse the sign of ΔH. When you multiply by a coefficient, multiply ΔH too. Keep track carefully as you add equations. Cancel species that appear on both sides.
Study Tips for Unit 6
Thermochemistry is packed with formulas and calculations. Master the essentials:
- Memorize key formulas early. q = mcΔT, q = n × ΔH, ΔH° = Σ(ΔH°f products) - Σ(ΔH°f reactants). These appear over and over.
- Practice sign conventions. Negative ΔH is exothermic. Positive is endothermic. Heat released by reaction is negative (leaves the system).
- Work through calorimetry problems step by step. Calculate heat from water, use energy conservation to find reaction enthalpy. Organization prevents mistakes.
- Get comfortable with Hess's Law. Write out each equation clearly. Track ΔH separately as you add equations. Double-check that species cancel correctly.
- Draw energy diagrams. Visual representation helps you spot whether processes are exothermic or endothermic and estimate relative energy differences.
- Use formation enthalpy tables correctly. Remember that elements in their standard states have ΔH°f = 0. Different states (s, l, g) have different values.
Summary, Review Questions & Practice
You've covered all the topics in Unit 6: from endothermic and exothermic processes to calorimetry, phase changes, and calculating reaction enthalpies using multiple methods. Before you move on, test yourself with these scenario-based questions. If you can answer them confidently, you're ready for equilibrium and acids-bases.
Review Questions: Test Yourself
- In a calorimetry experiment, a 100 g sample of sodium metal is dissolved in 1000 g of water. The temperature of the water rises from 20°C to 45°C. Calculate the heat released by the reaction. Given: specific heat of water = 4.18 J/g°C. Is this an exothermic or endothermic process?
- Using the following data: A → B, ΔH = 50 kJ; B → C, ΔH = -30 kJ. Calculate ΔH for A → C. Is the overall process exothermic or endothermic?
- The standard enthalpy of combustion of methane is ΔH° = -890 kJ/mol. Write the thermochemical equation for the combustion of methane and explain what the negative sign means.
- A reaction absorbs 250 kJ to melt 5 moles of ice. Calculate the molar enthalpy of fusion (ΔHfus). Would you expect the reverse process (freezing) to have the same or opposite sign for ΔH?
- Using bond enthalpies, estimate ΔH for the reaction H2 + Cl2 → 2 HCl. Given: H-H = 436 kJ/mol, Cl-Cl = 243 kJ/mol, H-Cl = 432 kJ/mol.
Want more practice? Paste these questions into StarSpark to generate a full quiz with explanations.
Explore the Full AP Chemistry Study Guide
Unit 6 gives you the tools to predict whether reactions are favorable and calculate the energy involved. This foundation is essential for understanding equilibrium, spontaneity, and reaction feasibility.
Check out the full AP Chemistry study plan to see how this unit connects to the rest of the course.
Other Unit Reviews:
- AP Chemistry Unit 1: Atomic Structure and Properties
- AP Chemistry Unit 2: Compound Structure and Properties
- AP Chemistry Unit 3: Properties of Substances and Mixtures
- AP Chemistry Unit 4: Chemical Reactions
- AP Chemistry Unit 5: Kinetics
- AP Chemistry Unit 7: Equilibrium
- AP Chemistry Unit 8: Acids and Bases
- AP Chemistry Unit 9: Thermodynamics and Electrochemistry
For official AP Chemistry resources, visit apcentral.collegeboard.org.
This review is aligned with the AP Chemistry Course and Exam Description. AP is a registered trademark of the College Board, which was not involved in the production of this guide.
