Chemical change
Start with the big picture
The First Law expresses the change in a system’s internal energy as ΔU = q + w, with heat and work positive when energy enters the system. In chemical reactions, volume changes can produce pressure–volume work. Calorimetry connects reaction heat to energy changes: at constant volume, q_v = ΔU; at constant pressure, q_p = ΔH. Enthalpy, defined as H = U + PV, is therefore useful for reactions at constant pressure. Negative reaction enthalpy denotes an exothermic change, while positive reaction enthalpy denotes an endothermic one. Because ΔU and ΔH are state functions, reaction enthalpies can be combined using Hess’s law. The lesson also introduces standard enthalpy changes, heat capacity, gas-phase relationships, and cyclic or adiabatic processes.
What you'll learn
- Relate heat, work, and internal-energy change using the First Law.
- Distinguish constant-volume and constant-pressure calorimetry.
- Interpret reaction enthalpy signs for exothermic and endothermic changes.
- Explain how state functions support Hess’s law.
- Identify how heat capacity and gas changes relate to reaction energetics.
Continue your study
Work through the complete notes and reinforce the topic with the study tools available in the full lesson.
