Chapter 5 — Thermodynamics
Every chemical change moves energy around — as heat, as work, or locked into new bonds. Systems and state functions, the first law, enthalpy and calorimetry, Hess's law and the Born–Haber cycle, and finally entropy and Gibbs energy, which decide whether a reaction can happen at all.
A cold pack from a sports first-aid kit contains a small bag of water and some solid ammonium nitrate. Squeeze it, the bag bursts, the salt dissolves — and the pack becomes so cold that it numbs a sprained ankle. A self-heating can of coffee does the opposite: break the seal, calcium oxide meets water, and the drink is hot within minutes. Both changes happen by themselves, once started. One gives out heat and the other soaks it up. So "giving out energy" cannot be the whole reason why things happen. What, then, decides the direction of a chemical change?
Thermodynamics answers that question with a handful of quantities that depend only on the state of a system, not on its history. The chapter starts with the bookkeeping of energy: a system exchanges heat $q$ and work $w$ with its surroundings, and its internal energy changes by $\Delta U=q+w$ — the first law. At constant pressure the heat exchanged is the enthalpy change $\Delta H$, which we measure in calorimeters and combine with Hess's law to find enthalpies of formation, combustion, atomisation, bonds and crystal lattices (the Born–Haber cycle). Finally entropy $S$, a measure of how spread out energy and matter are, joins enthalpy in the Gibbs energy $G=H-TS$. The sign of $\Delta G$ predicts whether a process is spontaneous, and the standard value $\Delta_{r}G^{\circ}$ fixes the equilibrium constant of the next chapter. Thermodynamics is examined in every Board paper and carries heavy weight in JEE and NEET.
- classify systems as open, closed or isolated, distinguish state functions from path functions, and properties as extensive or intensive;
- state the first law of thermodynamics as $\Delta U=q+w$ with the IUPAC sign convention, and apply it to heating, compression and expansion;
- calculate pressure–volume work for irreversible expansion against a constant pressure, for reversible isothermal expansion of an ideal gas and for free expansion, and explain why the reversible work is the largest;
- define enthalpy, use $\Delta H=\Delta U+\Delta n_{g}RT$, work with heat capacities ($q=mc\Delta T$, $C_{p}-C_{v}=R$), and interpret bomb-calorimeter and coffee-cup calorimeter measurements;
- write thermochemical equations, use standard enthalpies of formation, combustion, atomisation, bonds, solution and dilution, and apply Hess's law, including the Born–Haber cycle for a lattice enthalpy;
- explain what makes a process spontaneous, calculate entropy changes, and use the second law;
- use $\Delta G=\Delta H-T\Delta S$ to predict spontaneity in the four sign cases, find the crossover temperature, and connect $\Delta_{r}G^{\circ}$ with the equilibrium constant through $\Delta_{r}G^{\circ}=-RT\ln K$.
Sections in this chapter
Systems, Work, Heat and the First Law
System and surroundings; open, closed and isolated systems; state functions and internal energy; heat, work and the sign convention; the first law; pressure–volume work, reversible and irreversible; extensive and intensive properties.
Enthalpy, Heat Capacity and Calorimetry
Enthalpy as the heat exchanged at constant pressure; $\Delta H$ and $\Delta U$ for reactions involving gases; heat capacity, specific and molar heat capacities, $C_{p}-C_{v}=R$; bomb and coffee-cup calorimeters.
Thermochemistry: Enthalpies of Reaction and Hess's Law
Standard enthalpies of reaction and thermochemical equations; phase changes; enthalpies of formation, combustion, atomisation and bonds; Hess's law; lattice enthalpy and the Born–Haber cycle; enthalpies of solution and dilution.
Entropy, Gibbs Energy and Spontaneity
Why spontaneous processes happen; entropy and $\Delta S=q_{rev}/T$; the second law; Gibbs energy and the four sign cases of $\Delta G=\Delta H-T\Delta S$; $\Delta_{r}G^{\circ}$ and the equilibrium constant; the third law.