AP Chemistry Unit 6 Cheatsheet: Thermochemistry

Exothermic: System loses energy → surroundings gain heat. ΔH < 0. Temperature of surroundings increases.

Endothermic: System gains energy → surroundings lose heat. ΔH > 0. Temperature of surroundings decreases.

Key observation: In a calorimeter, warming of the surroundings means the system released heat, while cooling means the system absorbed heat.

• Dissolution can be either (depends on relative strength of intermolecular forces before/after)
• Phase changes: melting/boiling absorb energy (endothermic); freezing/condensing release energy (exothermic)

Formula: q = mcΔT

• q = heat absorbed or released (J)
• m = mass (g)
• c = specific heat capacity (J/g·°C)
• ΔT = temperature change (°C or K)

Specific heat capacity: Energy needed to raise 1 g of substance by 1 °C.

Molar heat capacity: Energy needed to raise 1 mole by 1 °C.

Key principle: Same thermal energy produces different ΔT in substances with different c values.

First Law of Thermodynamics: Energy is conserved in all processes.

In calorimetry: Heat lost by system = Heat gained by surroundings (or vice versa).

Dissolution calorimetry:

• Temperature increases → Exothermic dissolution (ΔH < 0)
• Temperature decreases → Endothermic dissolution (ΔH > 0)

Three main energy processes in chemical systems:

1. Heating/cooling (q = mcΔT)
2. Phase transitions (q = n × ΔH_phase)
3. Chemical reactions (q = n × ΔH_rxn)