Thermodynamics

Thermodynamics studies the energy changes that accompany physical and chemical processes. The part of the universe under study is the system; everything else is the surroundings. A system is open (exchanges matter and energy), closed (energy only) or isolated (neither) — a distinction NEET often opens with.

The total energy stored in a system is its internal energy $U$. We cannot measure $U$ itself, only its change $\Delta U$. The first law of thermodynamics is simply the conservation of energy: $\Delta U = q + W$, where $q$ is heat added to the system and $W$ is work done on it. For a gas expanding or compressing against an external pressure, the work is $W = -P\Delta V$ (work done by the gas is negative for the system).

A vital idea is the difference between state functions and path functions. A state function depends only on the present state, not on how it was reached — $U$, enthalpy $H$, entropy $S$ and Gibbs energy $G$ are all state functions. Heat $q$ and work $W$ are path functions: their values depend on the route taken. Knowing which quantities are state functions is a recurring NEET point.

Sign conventions matter. Heat absorbed by the system is positive, heat released is negative; work done on the system is positive, work done by the system is negative. For special processes the first law simplifies: at constant volume $W = 0$ so $\Delta U = q_V$, and in an isothermal process for an ideal gas $\Delta U = 0$ (internal energy depends only on temperature). Getting these signs right turns most numerical questions into one-line substitutions.

Most chemistry happens at constant (atmospheric) pressure, where it is convenient to define enthalpy $H = U + PV$. At constant pressure the heat exchanged equals the enthalpy change: $q_p = \Delta H$. For reactions involving gases, enthalpy and internal energy are related by $\Delta H = \Delta U + \Delta n_g RT$, where $\Delta n_g$ is the change in the number of gas moles — a formula NEET tests directly.

The sign of $\Delta H$ classifies reactions: exothermic reactions release heat ($\Delta H < 0$, e.g. combustion) and endothermic reactions absorb it ($\Delta H > 0$, e.g. photosynthesis, decomposition). Several standard enthalpies are defined: of formation (forming one mole of a compound from its elements), combustion, neutralisation, and bond enthalpy.

Hess's law of constant heat summation is the most powerful tool here. Because enthalpy is a state function, the total enthalpy change of a reaction is the same whether it occurs in one step or several. This lets you add the enthalpies of individual steps (reversing a step flips the sign; multiplying a step multiplies its $\Delta H$) to find an enthalpy that is hard to measure directly.

Hess's law also gives the standard enthalpy of reaction from formation values: $\Delta H_{rxn} = \sum \Delta H_f(\text{products}) - \sum \Delta H_f(\text{reactants})$, remembering that the $\Delta H_f$ of an element in its standard state is zero. Reaction enthalpy can likewise be estimated from bond enthalpies as (bonds broken) − (bonds formed). These three routes — steps, formation values, and bond enthalpies — cover essentially every NEET enthalpy calculation.

The first law tells us energy is conserved, but it does not say which way a process will go. That direction is governed by entropy $S$, a measure of the disorder or number of ways the energy and particles of a system can be arranged. The more disordered a state, the higher its entropy, and entropy — like enthalpy — is a state function.

Entropy increases with disorder, so for the same substance $S_{gas} > S_{liquid} > S_{solid}$. It rises on melting, boiling and dissolving, when a solid decomposes into gases, or when the number of gas molecules increases in a reaction. NEET frequently asks you to predict the sign of $\Delta S$ from such changes, so reading the disorder of reactants versus products is the key skill.

The second law of thermodynamics states that the total entropy of the universe always increases for a spontaneous process: $\Delta S_{universe} = \Delta S_{system} + \Delta S_{surroundings} > 0$. A process can lower the entropy of the system (for instance, freezing) only if the surroundings gain even more entropy, keeping the universe's total rising.

The third law adds a reference point: the entropy of a perfect crystalline solid is zero at absolute zero ($0\ \text{K}$), because there is only one way to arrange a perfectly ordered, motionless crystal. This allows absolute entropies to be tabulated. For NEET the practical takeaways are: predict the sign of $\Delta S$ from changes in state and gas moles, and remember that a spontaneous change always raises the entropy of the universe.

Tracking the entropy of the universe is awkward, so chemists use a property of the system alone that still predicts spontaneity: the Gibbs free energy $G = H - TS$. For a process at constant temperature and pressure, the change is $\Delta G = \Delta H - T\Delta S$. This single equation is the heart of chemical thermodynamics for NEET.

The sign of $\Delta G$ decides the direction: a process is spontaneous (feasible) when $\Delta G < 0$, non-spontaneous when $\Delta G > 0$, and at equilibrium when $\Delta G = 0$. The magnitude of $\Delta G$ also represents the maximum useful (non-expansion) work a reaction can deliver.

Because $\Delta G$ combines enthalpy and entropy with temperature, four cases arise. If $\Delta H < 0$ and $\Delta S > 0$, $\Delta G$ is negative at all temperatures (always spontaneous). If $\Delta H > 0$ and $\Delta S < 0$, it is positive at all temperatures (never spontaneous). The two mixed cases are temperature-dependent: an exothermic reaction with $\Delta S < 0$ is spontaneous only at low temperature, while an endothermic reaction with $\Delta S > 0$ becomes spontaneous only at high temperature. This explains why an endothermic reaction can still proceed.

Gibbs energy also links to equilibrium. The standard free-energy change relates to the equilibrium constant by $\Delta G^{\circ} = -RT\ln K$, so a large negative $\Delta G^{\circ}$ means a large $K$ (products favoured). Mastering $\Delta G = \Delta H - T\Delta S$, the four spontaneity cases, and $\Delta G^{\circ} = -RT\ln K$ covers nearly every spontaneity question NEET asks.