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🔥 Unit 6 · Thermodynamics 🏠 Unit Hub 🗂 Flashcards 🗺 Cheat Sheet ⭐ Essentials 🎨 Visual Review 📝 MC Practice ✍️ SAQ Practice

AP Chemistry Unit 6 Essentials

The must-know terms and big ideas for Unit 6: Thermodynamics. Every vocabulary word and concept you need to master.

← Back to Unit 6 hub
Big Idea 1
Energy always flows — the question is which direction
Every physical and chemical process either absorbs energy from its surroundings (endothermic) or releases energy into its surroundings (exothermic). Heat always flows spontaneously from a hotter object to a colder one until both reach thermal equilibrium — this single rule underlies calorimetry, phase changes, and every energy diagram you'll draw in this unit. Phase changes absorb or release energy without changing temperature, because that energy goes entirely into overcoming (or forming) intermolecular forces, not increasing kinetic energy.
Endothermic/Exothermic Heat Transfer Phase Changes
Big Idea 2
Enthalpy is a state function — so there are multiple paths to the same answer
Because enthalpy depends only on initial and final states (not the path taken to get there), chemists have three independent ways to calculate a reaction's enthalpy change: estimating it from bond enthalpies (energy to break reactant bonds minus energy to form product bonds), calculating it from tabulated standard enthalpies of formation, or combining known reactions algebraically using Hess's law. All three methods, applied correctly to the same reaction, must give the same ΔH — and recognizing this gives you flexibility on the exam to use whichever data you're given.
Enthalpy Bond Enthalpy Hess's Law
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Endothermic process
A process that absorbs energy from the surroundings, resulting in a positive ΔH (or ΔE) and often a temperature decrease in the surroundings.
Energy Changes
Exothermic process
A process that releases energy into the surroundings, resulting in a negative ΔH (or ΔE) and often a temperature increase in the surroundings.
Energy Changes
System vs. surroundings
The system is the part of the universe being studied (e.g., the reacting chemicals); the surroundings are everything else that can exchange energy with the system.
Energy Changes
Heat (q)
Energy transferred between a system and its surroundings due to a temperature difference; distinct from temperature itself.
Heat Transfer
Thermal equilibrium
The state reached when two objects in contact reach the same temperature and net heat flow between them stops.
Heat Transfer
Specific heat capacity (c)
The amount of heat required to raise the temperature of 1 gram of a substance by 1°C; a measure of a substance's resistance to temperature change.
Calorimetry
q = mcΔT
The calorimetry equation relating heat transferred (q) to mass (m), specific heat capacity (c), and temperature change (ΔT).
Calorimetry
Calorimetry
An experimental technique for measuring the heat absorbed or released during a physical or chemical process, often using a calorimeter to insulate the system.
Calorimetry
Heat of fusion / heat of vaporization
The energy required to melt a solid (fusion) or vaporize a liquid (vaporization) at constant temperature, without changing kinetic energy — the energy goes entirely into overcoming intermolecular forces.
Phase Changes
Enthalpy (H)
A state function representing the total heat content of a system at constant pressure; only changes in enthalpy (ΔH), not absolute values, are typically measured.
Enthalpy
State function
A property that depends only on the current state of a system (initial and final values), not on the path taken to reach that state. Enthalpy is a state function.
Enthalpy
Enthalpy of reaction (ΔHrxn)
The heat released or absorbed during a chemical reaction at constant pressure; negative for exothermic reactions, positive for endothermic reactions.
Enthalpy
Bond enthalpy
The energy required to break one mole of a particular covalent bond in the gas phase; used to estimate ΔHrxn as (bonds broken) − (bonds formed).
Bond Enthalpy
Standard enthalpy of formation (ΔHf°)
The enthalpy change when one mole of a compound forms from its elements in their standard states; by definition, ΔHf° of an element in its standard state is zero.
Enthalpy of Formation
Hess's law
States that the enthalpy change of an overall reaction equals the sum of the enthalpy changes of any set of steps that add up to that reaction, since enthalpy is a state function.
Hess's Law