Internal energy and enthalpy
Start with the big picture
Internal energy, U, represents a system’s microscopic kinetic and potential energy and depends on its state. The First Law expresses its change as ΔU = q + w, with heat absorbed and work done on the system taken as positive under the chemistry convention. For pressure–volume work, expansion gives a negative work term; at constant volume, ΔU equals the heat exchanged. Enthalpy is defined as H = U + PV, and at constant pressure its change equals the heat exchanged. The topic also introduces ideal-gas relations between ΔH and ΔU, heat capacities, reaction enthalpy, and the distinction between extensive and intensive properties. Together, these ideas help identify what can be inferred from a process’s conditions and which quantities depend on its path.
What you'll learn
- Apply the First Law using the chemistry sign convention for heat and work.
- Distinguish state functions from path functions in thermodynamic processes.
- Relate internal energy and enthalpy changes to constant-volume and constant-pressure heat.
- Interpret the roles of PV work and ideal-gas relations in energy changes.
- Identify how heat capacities and reaction enthalpy describe thermodynamic changes.
Continue your study
Work through the complete notes and reinforce the topic with the study tools available in the full lesson.
