Thermodynamics Notes for NEET and JEE Main 2027: Key Points and PDF

Thermodynamics is Unit 5 of NCERT Class 11 Chemistry (Part I) and the Chemical Thermodynamics unit of both the NEET UG and JEE Main syllabi. These notes cover system types, state functions, the first law, work, enthalpy, heat capacity, calorimetry, Hess’s law, types of enthalpy change, entropy, Gibbs energy and equilibrium, with a free notes PDF.

Thermodynamics Notes for NEET and JEE Main

Section-wise key points from NCERT Unit 5, formula tables, high-yield facts and common mistakes for NEET and JEE Main 2027 revision.

Thermodynamics at a Glance

The topic list follows the NEET UG 2026 syllabus released by the National Medical Commission (NMC) on 22 December 2025 and NTA’s JEE Main 2026 syllabus, which stay the reference until the 2027 syllabi are out. The full NEET list is on the NEET syllabus page.

System, Surroundings and State Functions

The system is the part of the universe under observation, the surroundings are everything else, and a real or imaginary boundary separates the two.

• The state of a system is described by state variables such as pressure, volume, temperature and amount.
• A state function depends only on the initial and final states, not on the path: U, H, S, G, p, V and T are state functions.
• Heat (q) and work (w) depend on the path, so they are not state functions.
• Internal energy (U) is the total energy of the system. It changes when heat passes into or out of the system, when work is done on or by it, or when matter enters or leaves.
• An adiabatic process allows no heat exchange between system and surroundings (q = 0); the wall that prevents it is an adiabatic wall.

What does the first law of thermodynamics say?

The first law says the energy of an isolated system is constant: energy can be neither created nor destroyed. In equation form, ΔU = q + w.

• NCERT sign convention: q and w are positive when heat or work is added to the system, and negative when the system loses them.
• Work done on a system in an adiabatic process raises its internal energy: ΔU = w(ad).
• Heat absorbed with no work done (constant volume) changes internal energy directly: ΔU = q(V).
• ΔU depends only on the initial and final states, even though q and w individually depend on the path.

How is pressure-volume work calculated?

Expansion work is w = –p(ex)ΔV, so work is negative when a gas expands (the system does work) and positive when it is compressed.

• A reversible process proceeds infinitely slowly through a series of equilibrium states and can be reversed at any moment by an infinitesimal change.
• Processes other than reversible ones are irreversible.

What is enthalpy and how is it related to internal energy?

Enthalpy is H = U + pV. At constant pressure the heat absorbed or released equals the enthalpy change: ΔH = q(p).

• ΔH = ΔU + pΔV. For reactions involving gases, ΔH = ΔU + Δn(g)RT.
• Δn(g) = moles of gaseous products minus moles of gaseous reactants. Liquids and solids are not counted.
• When Δn(g) = 0, or only solids and liquids take part, ΔH and ΔU are nearly equal.
• ΔH is negative for exothermic reactions (heat released) and positive for endothermic reactions (heat absorbed).
• Heat capacity: q = CΔT. With specific heat c and mass m, q = c × m × ΔT. Molar heat capacity C(m) is the heat capacity of one mole.
• For an ideal gas, C(p) – C(V) = R (per mole).

How are ΔU and ΔH measured?

Both are measured by calorimetry: ΔU at constant volume in a bomb calorimeter and ΔH at constant pressure.

• For an exothermic reaction at constant pressure, heat is evolved, so q(p) and Δ(r)H are negative.

Standard Enthalpy, Enthalpy of Formation and Hess’s Law

Enthalpy changes are compared in standard states: the standard state of a substance at a given temperature is its pure form at 1 bar.

• Standard enthalpy of reaction, Δ(r)H°, is the enthalpy change when all reactants and products are in their standard states. Data are usually taken at 298 K.
• Enthalpies of fusion, vaporisation and sublimation are phase-transformation enthalpies; they are positive, and melting and boiling take place at constant temperature.
• Standard molar enthalpy of formation, Δ(f)H°, is the enthalpy change when one mole of a compound forms from its elements in their most stable states (reference states).
• Δ(f)H° of an element in its reference state is zero. Reference states: H2(g), O2(g), C(graphite), S(rhombic).
• Δ(r)H° = Σ a(i) Δ(f)H°(products) – Σ b(i) Δ(f)H°(reactants), where a and b are the coefficients in the balanced equation.
• A thermochemical equation is a balanced equation with physical states and Δ(r)H. Reversing it changes the sign of Δ(r)H; multiplying it multiplies Δ(r)H.
• Hess’s law of constant heat summation: if a reaction takes place in several steps, its standard reaction enthalpy is the sum of the standard enthalpies of the intermediate steps at the same temperature.

What are the types of enthalpy change?

Each type is named after the process it describes and is given per mole.

• Gas-phase reactions: Δ(r)H° = Σ bond enthalpies of reactants – Σ bond enthalpies of products.
• The Born-Haber cycle applies Hess’s law: the sum of enthalpy changes round a complete cycle is zero.

What decides whether a process is spontaneous?

Neither a decrease in enthalpy nor an increase in the system’s entropy alone decides spontaneity. The total entropy change, or at constant temperature and pressure the Gibbs energy change, does.

Entropy

• Entropy (S) is a measure of the degree of randomness or disorder in a system; it is a state function.
• For a given substance the crystalline solid has the lowest entropy and the gas the highest.
• ΔS = q(rev)/T. The same heat causes more randomness when added at a lower temperature.
• For a spontaneous process, ΔS(total) = ΔS(system) + ΔS(surroundings) > 0. At equilibrium entropy is at its maximum and ΔS = 0.
• For an isolated system ΔU = 0, yet ΔS > 0 for a spontaneous change, so entropy (not energy) identifies the change.
• ΔS(surroundings) = –ΔH(system)/T. NCERT example: rusting of iron has ΔS = –549.4 J/K/mol, but ΔS(surroundings) is about +5530 J/K/mol, so the total is positive and the reaction is spontaneous.
• Second law: the entropy of an isolated system increases in a spontaneous change.
• Third law: the entropy of a pure crystalline substance approaches zero as the temperature approaches absolute zero. It allows absolute entropies to be calculated from thermal data.

Gibbs Energy

• G = H – TS. Gibbs energy is an extensive property and a state function.
• At constant temperature, ΔG = ΔH – TΔS (the Gibbs equation).
• ΔG < 0: spontaneous. ΔG > 0: non-spontaneous. ΔG = 0: equilibrium.
• ΔG is the net energy available to do useful work, so it is also called free energy; TΔS is the energy not available for useful work.

How is Gibbs energy related to the equilibrium constant?

At equilibrium Δ(r)G = 0, and the standard Gibbs energy change fixes the equilibrium constant: Δ(r)G° = –RT ln K = –2.303 RT log K.

• Δ(r)G° = Δ(r)H° – TΔ(r)S° = –RT ln K, so K can be calculated from enthalpy and entropy data.
• A large negative Δ(r)G° (for example, a strongly exothermic reaction) gives K much greater than 1, so the reaction goes nearly to completion.
• A large positive Δ(r)G° (for example, a strongly endothermic reaction) gives K much smaller than 1.
• Temperature matters: reactions with positive Δ(r)S that are non-spontaneous at low temperature can become spontaneous at high temperature.

Most-asked Facts for NEET and JEE Main: Thermodynamics

Revise these one-line facts from the NCERT chapter just before the exam.

• First law: ΔU = q + w, with q and w positive when added to the system.
• Open system: energy and matter exchanged; closed: energy only; isolated: neither.
• U, H, S and G are state functions; q and w are path functions.
• w = –p(ex)ΔV; isothermal reversible work = –2.303 nRT log (V(f)/V(i)).
• Free expansion of an ideal gas: w = 0.
• Isothermal ideal gas: ΔU = 0 and q = –w. Adiabatic process: q = 0.
• ΔU = q(V) (bomb calorimeter); ΔH = q(p).
• ΔH = ΔU + Δn(g)RT, counting only gaseous species.
• C(p) – C(V) = R for one mole of an ideal gas.
• Standard state: pure substance at 1 bar.
• Δ(f)H° of an element in its reference state, such as graphite or rhombic sulphur, is zero.
• Hess’s law: total enthalpy change is the same whatever the path.
• Mean C–H bond enthalpy in methane is about 416 kJ/mol (1665/4).
• Lattice enthalpy of NaCl is +788 kJ/mol; Δ(sol)H = Δ(lattice)H + Δ(hyd)H.
• ΔS = q(rev)/T; spontaneous change has ΔS(total) > 0.
• Third law: entropy of a pure crystal tends to zero at 0 K.
• ΔG = ΔH – TΔS; ΔG < 0 means spontaneous at constant T and p.
• Δ(r)G° = –2.303 RT log K.

Common Mistakes to Avoid

Watch for these mix-ups when you answer questions on this chapter.

• Using the Physics sign convention in Chemistry. NCERT Chemistry writes ΔU = q + w, with work done on the system positive; NCERT Physics counts work done by the system as positive.
• Counting liquids or solids in Δn(g). Only gaseous reactants and products count.
• Mixing units in ΔG = ΔH – TΔS. ΔS is usually in J/K/mol and ΔH in kJ/mol; convert before subtracting.
• Writing products minus reactants for bond enthalpies. For bond enthalpies it is reactants minus products, the reverse of the enthalpy-of-formation formula.
• Taking Δ(f)H° = 0 for any form of an element. Only the reference state is zero: graphite, not diamond.
• Forgetting to change the sign of ΔH when a thermochemical equation is reversed.
• Assuming every exothermic reaction is spontaneous. A reaction with negative ΔH and negative ΔS becomes non-spontaneous at high temperature.
• Using ΔS = q/T for an irreversible path. The relation uses the heat of a reversible process.
• Using the gas pressure instead of the external pressure for irreversible work.

Thermodynamics Notes PDF and Related Pages

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• Source text: NCERT Unit 5 (NCERT, Class 11 Chemistry Part I)

Chemistry Notes for NEET and JEE Main: All Chapters

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• Structure of Atom
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