AP Chemistry Unit 6: Thermochemistry Notes

Introduction

Thermochemistry tracks energy across a declared system boundary. Heat signs describe transfer relative to the system; temperature is a state measure, not an amount of heat stored in an object.

Calorimetry and phase-change calculations apply energy conservation with explicit masses, heat capacities, states, and signs. A measured temperature change must be placed inside a complete system-plus-surroundings balance.

Enthalpy is a state function. Reaction scaling, bond-energy estimates, formation tables, and Hess combinations use different data paths but must preserve physical states, coefficients, and product-minus-reactant sign conventions.

Current exam guidance: Aligned to the current 2026–27 AP Chemistry course framework and exam scope.

6.1

Endothermic and Exothermic Processes

How does system choice determine whether a process is endothermic or exothermic?

Short answer

Heat entering the declared system is positive and endothermic; heat leaving it is negative and exothermic, with the surroundings carrying the equal opposite transfer.

Know firstSystem and surroundings · Signed quantities · Energy conservation

Name the system before assigning a sign

Endothermic and exothermic describe heat transfer relative to a chosen system. If the system absorbs heat, $q_{sys}>0$ ; if it releases heat, $q_{sys}<0$ . The same transfer has the opposite sign for the surroundings.

At constant pressure, heat transferred to the system equals its enthalpy change under the stated conditions. A thermometer often measures surroundings or solution temperature, so a temperature rise in the surroundings can indicate an exothermic system process.

Bond breaking or separating attractions requires energy; bond formation or stronger new attractions releases it. Dissolution can therefore be endothermic or exothermic depending on the balance among separated solute/solvent interactions and newly formed solute–solvent interactions. Overall sign comes from the complete before-and-after balance.

Chemistry model

Heat direction depends on the chosen system

Place system and surroundings boundaries before assigning signs to heat flow.

Surroundings

Systemreaction mixture

qsystem < 0

If heat enters system: qsystem > 0
If heat leaves system: qsystem < 0
Energy ledger: qsystem = −qsurroundings

q_{\mathrm{system}}=-q_{\mathrm{surroundings}}

Read this figure without the graphic

Heat direction depends on the chosen system. The figure connects system, surroundings, heat direction so the relationship can be read without relying on color alone.

Representation	What to read
system	starting evidence or quantity
surroundings	intermediate relationship
heat direction	result or conclusion

Formula

Paired heat signs

q_{sys}=-q_{surr},\qquad q_p=\Delta H_{sys}

Energy lost by one part is gained by the other; constant-pressure system heat uses the enthalpy sign.

Use when: The combined system and surroundings are treated as energy-conserving for the transfer and pressure is constant for $q_p=\Delta H$ .
Do not assume: A temperature rise in one object does not by itself label the chosen chemical system.
Variables: $q$ : heat transfer (J or kJ)
Common misuse: Saying bond breaking releases energy.

Signed transfer ledger

Record	Heat change	Classification relative to that record
Surroundings	−2.50 kJ	loses heat
System	+2.50 kJ	endothermic
Combined	0.00 kJ	energy conserved

Keep these three ideas

Heat sign is relative to the declared system.
System and surroundings heat changes are equal and opposite.
Bond breaking absorbs energy; full balance determines reaction sign.

Thermochemistry