Board exam focus — Chemical Thermodynamics (CBSE & HBSE)

Thermodynamics studies energy changes accompanying physical and chemical processes — without worrying about how fast or by what path they occur. This chapter introduces systems, surroundings and state functions; the first law (energy conservation), internal energy, work and heat; enthalpy and heat capacity; thermochemistry (enthalpies of reaction, formation, combustion and Hess's law); and finally spontaneity, entropy, the second law and Gibbs free energy ΔG = ΔH − TΔS. It is numerically rich and a strong scoring chapter.

System, Surroundings, State Functions, First Law & Enthalpy

Basic Terms

• System — the part of the universe under study.
• Surroundings — everything else; Universe = System + Surroundings.
• Boundary — the wall separating system and surroundings.

State Functions vs Path Functions

• State function — depends only on initial and final states, not the path. Examples: internal energy (U), enthalpy (H), entropy (S), Gibbs energy (G), T, P, V.
• Path function — depends on the path taken. Examples: heat (q) and work (w).

Trap: q and w are NOT state functions individually, but their algebraic sum (q + w = ΔU) IS a state function.

Internal Energy (U)

Total energy (kinetic + potential) stored in a system. We measure only changes (ΔU), not absolute values.

Sign Conventions (IUPAC)

• Heat absorbed BY system: q is +; released: q is −.
• Work done ON system: w is +; work done BY system: w is −.

First Law of Thermodynamics

Energy can neither be created nor destroyed — the energy of an isolated universe is constant.

$$\Delta U = q + w$$

For pressure-volume work at constant external pressure: w = −Pₑₓₜ ΔV. So ΔU = q − Pₑₓₜ ΔV.

• At constant volume (ΔV = 0): ΔU = qᵥ (heat measured in a bomb calorimeter).

Enthalpy (H)

Defined as H = U + PV. It is a state function. At constant pressure:

$$\Delta H = q_p$$

For reactions involving gases: ΔH = ΔU + Δnₘ RT, where Δnₘ = (moles of gaseous products − moles of gaseous reactants).

• Exothermic: ΔH < 0 (heat released).
• Endothermic: ΔH > 0 (heat absorbed).

Heat Capacity

• Cₚ (at constant pressure) and Cᵥ (at constant volume).
• For an ideal gas: Cₚ − Cᵥ = R (Mayer's relation).
• Heat supplied: q = m·c·ΔT (c = specific heat) or q = n·C·ΔT.

Thermochemistry: Enthalpies of Reaction, Formation, Combustion & Hess's Law

Standard Enthalpy of Reaction (ΔᵣH°)

The enthalpy change when reactants in their standard states (1 bar, specified T usually 298 K) convert to products in their standard states.

$$\Delta_r H^\circ = \sum \Delta_f H^\circ (\text{products}) - \sum \Delta_f H^\circ (\text{reactants})$$

Types of Enthalpy Changes

Hess's Law of Constant Heat Summation

The total enthalpy change of a reaction is the same whether it occurs in one step or several steps, provided initial and final conditions are identical.

This follows directly from H being a state function. It lets us add thermochemical equations algebraically (and reverse them, flipping the sign).

Enthalpy from Bond Enthalpies

For gas-phase reactions: $$\Delta_r H = \sum (\text{bond enthalpies of reactants}) - \sum (\text{bond enthalpies of products})$$

Trap: This formula is bonds broken − bonds formed. Be careful with the direction — getting the sign backwards is the most common error.

Rules for thermochemical equations: reversing a reaction flips the sign of ΔH; multiplying coefficients multiplies ΔH by the same factor.

Spontaneity, Entropy, Second Law & Gibbs Free Energy

Spontaneous Process

A process that occurs on its own (without external help), tending toward equilibrium. Spontaneity is about feasibility, not speed.

Many spontaneous processes are exothermic, but NOT all — melting of ice and dissolution of NH₄Cl are spontaneous yet endothermic. So ΔH alone cannot decide spontaneity; entropy matters too.

Entropy (S)

A state function measuring the degree of disorder/randomness of a system.

• Entropy increases: solid → liquid → gas, dissolution, increase in number of gas moles, rise in temperature.
• ΔS = qᵣₑᵥ / T (for a reversible process at temperature T).

Second Law of Thermodynamics

For any spontaneous process, the total entropy of the universe increases:

$$\Delta S_{total} = \Delta S_{system} + \Delta S_{surroundings} > 0$$

At equilibrium, ΔS_total = 0.

Gibbs Free Energy (G)

Defined as G = H − TS. At constant T and P:

$$\Delta G = \Delta H - T\Delta S$$

ΔG is the criterion of spontaneity:

Effect of ΔH, ΔS and Temperature on ΔG

Gibbs Energy and Equilibrium Constant

$$\Delta G^\circ = -RT \ln K = -2.303\, RT \log K$$

• K > 1 ⇒ ΔG° < 0 (products favoured).
• The temperature at which a reaction changes from non-spontaneous to spontaneous: T = ΔH / ΔS (where ΔG = 0).

Third Law of Thermodynamics

The entropy of a perfectly crystalline substance is zero at absolute zero (0 K).

Frequently asked questions

Are these Chemical Thermodynamics notes free?

Yes — the Chemical Thermodynamics notes for Chemistry (Class 11) on Siksha Sarovar are completely free to read, with no account required.

Do these notes follow CBSE and HBSE?

Yes. The Chemical Thermodynamics notes are NCERT-aligned and include guidance for both CBSE and Haryana Board (HBSE), with important questions and MCQs for revision.

What does the Chemical Thermodynamics chapter cover?

Concept explanations, key formulas and definitions, fully solved examples and board-pattern practice questions for Chemical Thermodynamics.