AP Chemistry Unit 6 Thermodynamics: Understanding Enthalpy and Heat Changes in Reactions

Enthalpy (H)

A thermodynamic quantity used to track energy changes at constant pressure; defined as H = U + PV.

Internal energy (U)

The energy associated with the particles in a system (microscopic kinetic + potential energy), used in the definition H = U + PV.

PV term

The “pressure–volume” energy correction in enthalpy that accounts for energy involved in pushing back the atmosphere as volume changes.

Enthalpy change of reaction (ΔHrxn)

The difference in enthalpy between products and reactants: ΔHrxn = Hproducts − Hreactants.

Exothermic reaction

A reaction with ΔHrxn < 0 that releases heat to the surroundings (surroundings warm as the system loses heat).

Endothermic reaction

A reaction with ΔHrxn > 0 that absorbs heat from the surroundings (surroundings cool as the system gains heat).

Constant-pressure heat (qp)

Heat absorbed by the system at constant pressure; at constant pressure, qp = ΔH.

Thermochemical equation

A balanced chemical equation that includes an enthalpy change (ΔH) tied to the reaction exactly as written (including coefficients and states).

Extensive property (in reaction enthalpy context)

A property that scales with the amount of substance; if you scale a reaction by a factor, ΔH scales by the same factor.

Scaling ΔH with stoichiometry

Rule that ΔH corresponds to the balanced equation’s coefficients: doubling the reaction doubles ΔH; reversing the reaction changes the sign of ΔH.

Calorimetry

An experimental method to determine heat flow by measuring temperature change, then relating that heat to ΔH for the reaction.

Specific heat capacity (c)

Heat required to raise 1 g of a substance by 1°C (or 1 K); used in q = mcΔT.

Heat capacity (C) of a calorimeter

Heat required to raise the calorimeter’s temperature by 1°C (or 1 K); used in q = CΔT.

Temperature change (ΔT)

The change in temperature (Tf − Ti) used in calorimetry calculations to infer heat transfer.

Bond enthalpy (bond dissociation enthalpy)

The enthalpy change required to break a specific bond in the gas phase; bond breaking requires energy input.

Average bond enthalpy

A tabulated bond enthalpy that is averaged over many molecules containing that bond, so calculations using it provide estimates rather than exact values.

Bond breaking

The process of pulling bonded atoms apart; it is endothermic and contributes positive energy to ΔH.

Bond formation

The process of forming a bond between atoms; it is exothermic and releases energy.

Bond-enthalpy estimate for reaction enthalpy

An approximate method: ΔHrxn ≈ ΣDbroken − ΣDformed (add energy for bonds broken, subtract energy released for bonds formed).

Standard enthalpy of formation (ΔHf°)

The enthalpy change when 1 mole of a compound forms from its elements in their standard states under standard conditions.

Standard state

The most stable form of an element under standard conditions (commonly 1 bar and a specified temperature, often 298 K), e.g., O2(g), H2(g), C(s, graphite), Br2(l).

ΔHf° of an element in its standard state

Defined as 0 kJ/mol (a reference point for the formation enthalpy scale, not “no enthalpy”).

Standard enthalpy of reaction from formation enthalpies (ΔHrxn°)

Calculated by ΔHrxn° = ΣνΔHf°(products) − ΣνΔHf°(reactants), where ν are stoichiometric coefficients.

State function

A property that depends only on the initial and final states, not on the path; enthalpy is a state function.

Hess’s Law

The enthalpy change of an overall reaction equals the sum of enthalpy changes of steps that add up to that overall reaction (because enthalpy is a state function).