AP Chemistry Thermochemistry (4 notes packets updated 12/20/2025)

Thermochemistry

The study of energy changes (heat) associated with chemical reactions and physical changes

Energy

The ability to do work or produce heat

Heat (q)

Energy transferred between a system and surroundings due to a temperature difference

System

The part of the universe being studied

Surroundings

Everything outside the system

Exothermic process

A process where heat is released from the system to the surroundings

q < 0

Endothermic process

A process where heat is absorbed by the system from the surroundings

q > 0

Heat of reaction (q_rxn)

The heat absorbed or released during a chemical reaction

Enthalpy (H)

The total heat content of a system

Enthalpy change (ΔH)

The heat of reaction at constant pressure

ΔH = q_p

At constant pressure the enthalpy change equals the heat transferred

Bond breaking

Requires energy (endothermic)

Bond forming

Releases energy (exothermic)

Overall ΔH sign

Depends on whether more energy is absorbed breaking bonds or released forming bonds

Specific heat capacity (c)

The amount of heat required to raise the temperature of 1 gram of a substance by 1°C

Units of specific heat

J/g·°C

Heat capacity (C)

The amount of heat required to raise the temperature of an entire object by 1°C

q = mcΔT

Equation used to calculate heat when mass and specific heat are known

q = CΔT

Equation used when heat capacity is given

ΔT

Final temperature minus initial temperature (Tf − Ti)

Calorimetry

The experimental measurement of heat flow

Calorimeter

A device used to measure heat absorbed or released during a reaction

Constant-pressure calorimeter

Measures heat changes at constant atmospheric pressure

qcal = −qrxn

Heat absorbed by the calorimeter equals the negative of the heat released by the reaction

Heat flow convention

Heat gained by surroundings is positive

heat lost by the system is negative

Thermochemical equation

A balanced chemical equation that includes the enthalpy change (ΔH)

ΔH applies to equation

ΔH is valid only for the reaction as written

Reversing a reaction

Reverses the sign of ΔH

Multiplying a reaction

Requires multiplying ΔH by the same factor

Adding reactions

Add their ΔH values

Stoichiometry with ΔH

Heat changes scale with the amount of substance reacting

Limiting reactant and heat

The limiting reactant determines the total heat released or absorbed

Hess’s Law

The enthalpy change of a reaction depends only on initial and final states not the path taken

Path independence

Different reaction pathways give the same overall ΔH

Using Hess’s Law

Reactions may be added reversed or multiplied to find unknown ΔH values

Phase change enthalpy

Energy changes associated with phase transitions

ΔH_fusion

Heat required to melt a solid

ΔH_vaporization

Heat required to vaporize a liquid

ΔH_sublimation

Heat required to convert a solid directly to a gas

ΔHsub = ΔHfus + ΔH_vap

Special relationship between phase change enthalpies

Standard heat of formation (ΔH°_f)

Enthalpy change when 1 mole of a compound forms from its elements in their standard states

Standard state

Pure substance at 1 atm and 25°C (298 K)

ΔH°_f of elements

Zero for elements in their standard states

Formation reaction rule

Formation reactions must produce exactly 1 mole of product

Calculating ΔH using ΔH°_f

ΔH°rxn = ΣΔH°f(products) − ΣΔH°_f(reactants)

Coefficients in ΔH° calculations

Each ΔH°_f must be multiplied by its balanced-equation coefficient

Physical state importance

ΔH°_f values depend on the phase of the substance

Methods to find heat

Calorimetry Hess’s Law and standard heats of formation

Units of enthalpy

kJ or kJ/mol

Sign of ΔH

Negative indicates exothermic positive indicates endothermic

Magnitude of ΔH

Directly proportional to the amount of substance reacting