Chemical Thermodynamics
Syllabus Unit: Physical Chemistry
Why it matters
Chapter overview
This topic quantifies the energy changes accompanying chemical reactions — the First Law's internal energy and enthalpy bookkeeping, Hess's Law for calculating unmeasurable enthalpies indirectly, the Second Law's entropy criterion for spontaneity, and Gibbs free energy, which unifies enthalpy and entropy into chemistry's single practical spontaneity test.
Chemistry Syllabus Physical Chemistry ~25% weightage
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1Heat Exchanged (q)Absorbed or released by system
2Work Done (w)On or by the system
3ΔU = q + wFirst Law bookkeeping
4ΔH = qpHeat term at constant pressure
1First Law & Enthalpy
The First Law says energy is conserved — a system's internal energy changes only by heat exchanged and work done; enthalpy is the convenient heat term at constant pressure.
Learning Objectives
- State the First Law of Thermodynamics: ΔU = q + w.
- Distinguish system/surroundings and open/closed/isolated systems.
- Define enthalpy H = U + PV and derive ΔH = ΔU + PΔV for constant-pressure processes.
- Classify processes as exothermic (ΔH negative) or endothermic (ΔH positive) with real examples.
Key Terms
Internal Energy (U)The total energy (kinetic + potential) stored within a system; a state function.
First Law of ThermodynamicsEnergy can neither be created nor destroyed, only converted between forms; ΔU = q + w.
Enthalpy (H)H = U + PV; the heat absorbed or released by a process at constant pressure, ΔH = qp.
State FunctionA property (like U, H, S, G) that depends only on the initial and final states, not the path taken to get there.
Thermodynamics is bookkeeping for heat and work exchanged between a system and its surroundings. The First Law is simply energy conservation applied to that exchange.
The First Law
ΔU = q + w (IUPAC convention: w = -PΔV for expansion work; work done ON the system is positive)
- q is positive when heat is absorbed by the system, negative when released.
- w is positive when work is done on the system.
- ΔU depends only on the initial and final states (a state function) — unlike q and w individually, which depend on the path taken.
Enthalpy
H = U + PV → ΔH = ΔU + PΔV (at constant pressure) → ΔH = qp
- Exothermic reaction — ΔH negative, heat released to the surroundings (e.g. combustion, neutralisation).
- Endothermic reaction — ΔH positive, heat absorbed from the surroundings (e.g. dissolving NH4NO3 in water, thermal decomposition of CaCO3).
Exam TipAt constant volume, qv = ΔU (no PV work is possible); at constant pressure, qp = ΔH. NEET numericals often hinge entirely on identifying which condition applies.
Common MistakeSign conventions for work differ between textbooks (w = +PΔV vs w = -PΔV) — always check which convention a numerical is using before substituting values.