Chemical Thermodynamics JEE Notes PDF, Download Now

Thermodynamics studies the energy changes associated with physical and chemical processes. Every time a reaction occurs, whether fuel burns or a salt dissolves, energy is released or absorbed, and thermodynamics explains how much energy changes, in which direction it flows and whether the process can occur spontaneously. Practising JEE questions on Chemical Thermodynamics helps students apply the first law, enthalpy, Hess's law, entropy and Gibbs free energy in numerical and conceptual problems. These Chemical Thermodynamics JEE notes cover the first law, enthalpy and standard enthalpy changes, Hess's law, the Born-Haber cycle, entropy, the second law and Gibbs free energy for quick JEE revision.

Chemical Thermodynamics JEE Notes: Important Concepts

Thermodynamics becomes much easier once the basic terminology is clear.

System: The specific part of the universe under study, such as a gas in a cylinder, a beaker of solution or a reaction mixture.

Surroundings: Everything outside the system that can exchange energy or matter with it.

Universe: The system together with its surroundings.

Types of Thermodynamic Systems

JEE tip: A thermos flask is the standard example used for an isolated system because ideally it exchanges neither matter nor energy with its surroundings.

State Functions and Path Functions

A state function depends only on the present state of a system and not on the path followed to reach it. Internal energy, enthalpy, entropy, pressure, temperature and volume are state functions.

A path function depends on the route followed during the process. Heat and work are path functions.

Heat and work are individually path functions, but their combined effect determines the change in internal energy, which is a state function.

Sign Conventions in Thermodynamics

• Heat absorbed by the system: Positive.
• Heat released by the system: Negative.
• Work done on the system: Positive.
• Work done by the system: Negative.

Internal Energy, First Law and Work

The total microscopic kinetic and potential energy stored in a system is called its internal energy. Its absolute value cannot normally be measured directly, so thermodynamics focuses on changes in internal energy.

First Law of Thermodynamics

The first law is the law of conservation of energy applied to a thermodynamic system.

$$\Delta U=q+w$$

For pressure-volume work against a constant external pressure:

$$w=-P_{\text{ext}}\Delta V$$

Therefore:

$$\Delta U=q-P_{\text{ext}}\Delta V$$

Worked example: A gas absorbs 500 J of heat and does 200 J of work on the surroundings. Heat absorbed is positive and work done by the system is negative.

$$\Delta U=500-200=300\text{ J}$$

Change in internal energy = +300 J.

Types of Thermodynamic Processes

Work Done by an Ideal Gas

For irreversible expansion or compression against a constant external pressure:

$$w=-P_{\text{ext}}(V_2-V_1)$$

For a reversible isothermal expansion or compression of an ideal gas:

$$w=-nRT\ln\left(\frac{V_2}{V_1}\right)$$

or:

$$w=-2.303nRT\log\left(\frac{V_2}{V_1}\right)$$

Worked example: Two moles of an ideal gas expand reversibly and isothermally at 300 K from 10 L to 20 L.

$$w=-(2)(8.314)(300)\ln(2)\approx-3.46\text{ kJ}$$

The negative sign shows that the gas does work on the surroundings.

Enthalpy and Standard Enthalpy Changes

Most chemical reactions are studied at constant atmospheric pressure. The heat exchanged under constant-pressure conditions is represented through enthalpy change.

Enthalpy

$$H=U+PV$$

At constant pressure:

$$\Delta H=q_p$$

For reactions involving ideal gases:

$$\Delta H=\Delta U+\Delta n_gRT$$

Here, $$\Delta n_g$$ is the number of moles of gaseous products minus the number of moles of gaseous reactants. Only gaseous species are counted.

Worked example: For $$N_2(g)+3H_2(g)\rightarrow2NH_3(g)$$, $$\Delta n_g=-2$$. If $$\Delta H=-92.4\text{ kJ}$$ at 298 K, then:

$$\Delta U\approx-87.4\text{ kJ}$$

Exothermic and Endothermic Reactions

Standard Enthalpy of Formation

The standard enthalpy of formation is the enthalpy change when one mole of a compound is formed from its elements in their standard states.

$$\Delta_rH^\circ=\sum\Delta_fH^\circ(\text{products})-\sum\Delta_fH^\circ(\text{reactants})$$

The standard enthalpy of formation of an element in its standard state is zero by definition.

Worked example: For methane combustion, $$CH_4(g)+2O_2(g)\rightarrow CO_2(g)+2H_2O(l)$$, the standard reaction enthalpy is:

$$\Delta_rH^\circ=-890.3\text{ kJ mol}^{-1}$$

Important Standard Enthalpy Changes

Bond Enthalpy

Breaking bonds requires energy, while forming bonds releases energy.

$$\Delta_rH^\circ=\sum(\text{bond energies of bonds broken})-\sum(\text{bond energies of bonds formed})$$

Worked example: For $$H_2(g)+Cl_2(g)\rightarrow2HCl(g)$$, bonds broken require 678 kJ and bonds formed release 862 kJ.

$$\Delta_rH^\circ=-184\text{ kJ mol}^{-1}$$

Hess's Law, Kirchhoff's Equation and Born-Haber Cycle

Hess's Law

Hess's law states that the total enthalpy change of a reaction is independent of the path followed because enthalpy is a state function.

$$\Delta H_{\text{reaction}}=\Delta H_1+\Delta H_2+\cdots$$

• Reversing a reaction reverses the sign of its enthalpy change.
• Multiplying a reaction by a factor multiplies its enthalpy change by the same factor.

Worked example: Combining the combustion of carbon with the reverse oxidation of carbon monoxide gives:

$$C+\frac12O_2\rightarrow CO$$

$$\Delta H=-110.5\text{ kJ}$$

Kirchhoff's Equation

Reaction enthalpy changes with temperature because reactants and products can have different heat capacities.

$$\Delta H_2=\Delta H_1+\Delta C_p(T_2-T_1)$$

where:

$$\Delta C_p=\sum C_p(\text{products})-\sum C_p(\text{reactants})$$

Worked example: If a reaction has an enthalpy change of −50 kJ at 300 K and $$\Delta C_p=-0.05\text{ kJ K}^{-1}$$, then at 500 K:

$$\Delta H_{500}=-60\text{ kJ}$$

Born-Haber Cycle

The Born-Haber cycle applies Hess's law to ionic solids and is commonly used to determine lattice enthalpy indirectly.

For NaCl, the cycle includes sublimation of sodium, ionisation of gaseous sodium, dissociation of chlorine, electron gain by chlorine and formation of the ionic lattice.

$$\Delta_fH^\circ=\Delta_{\text{sub}}H^\circ+\Delta_{\text{ion}}H^\circ+\frac12\Delta_{\text{bond}}H^\circ+\Delta_{\text{eg}}H^\circ-\Delta_{\text{lattice}}H^\circ$$

Worked example: Using the supplied NaCl data:

$$-411=108+496+121-349-\Delta_{\text{lattice}}H^\circ$$

$$\Delta_{\text{lattice}}H^\circ=787\text{ kJ mol}^{-1}$$

JEE tip: In Born-Haber cycle questions, write every step with its sign before substituting values. Most errors occur because a step is missed or a sign is reversed.

Entropy, Second Law and Gibbs Free Energy

Entropy

Entropy measures the degree of energy dispersal or disorder in a system. Gases generally have greater entropy than liquids, and liquids have greater entropy than solids.

$$S_{\text{gas}}>S_{\text{liquid}}>S_{\text{solid}}$$

For a reversible process:

$$\Delta S=\frac{q_{\text{rev}}}{T}$$

For a chemical reaction:

$$\Delta_rS^\circ=\sum S^\circ(\text{products})-\sum S^\circ(\text{reactants})$$

Second Law of Thermodynamics

For a spontaneous process, the total entropy of the universe increases.

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

Predicting Entropy Change

Entropy generally increases when a solid melts, a liquid vaporises, the number of gaseous molecules rises, temperature increases or a substance dissolves and becomes more dispersed.

Worked example: In $$CaCO_3(s)\rightarrow CaO(s)+CO_2(g)$$, a gas is produced from solid reactants, so disorder increases and the entropy change is positive.

Gibbs Free Energy

Gibbs free energy combines enthalpy and entropy to determine whether a process is thermodynamically spontaneous at constant temperature and pressure.

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

Effect of Enthalpy and Entropy on Spontaneity

When enthalpy change and entropy change have the same sign, a crossover temperature may exist.

$$T=\frac{\Delta H}{\Delta S}$$

Gibbs Free Energy and Equilibrium

The relationship between free energy and the reaction quotient is:

$$\Delta G=\Delta G^\circ+RT\ln Q$$

At equilibrium:

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

• If the standard free energy change is negative, products are favoured.
• If the standard free energy change is zero, neither side is favoured.
• If the standard free energy change is positive, reactants are favoured.

Chemical Thermodynamics Formula Sheet at a Glance

The first law, the enthalpy–internal energy relation and the Gibbs free energy relation are the three core equations that carry most of this chapter. Revising them regularly with a JEE formula sheet helps you recall sign conventions, standard thermodynamic relations and common numerical shortcuts more quickly.

JEE Important Points, Common Mistakes and Quick Revision

Points JEE Repeatedly Tests

• Internal energy, enthalpy and entropy are state functions, while heat and work are path functions.
• Heat absorbed by the system is positive, while work done by the system is negative under the chemistry sign convention.
• Only gaseous species are counted while calculating the change in gaseous moles.
• At constant pressure, heat exchanged equals enthalpy change.
• The standard enthalpy of formation of an element in its standard state is zero.
• Bond breaking requires energy, while bond formation releases energy.
• Hess's law works because enthalpy is a state function.
• Born-Haber cycles are commonly used to determine lattice enthalpy.
• Entropy generally increases when matter becomes more disordered.
• Spontaneity is determined by Gibbs free energy rather than enthalpy alone.
• A spontaneous reaction is not necessarily a fast reaction.

Common Mistakes to Avoid

1. Using the wrong sign for work. Expansion means the system does work, so work is negative under the IUPAC convention.
2. Counting solids and liquids in the change in gaseous moles. Only gaseous reactants and products are included.
3. Mixing joules and kilojoules. This is especially common in Gibbs free energy calculations.
4. Using the constant-pressure work expression for a reversible isothermal process. The reversible expression contains a logarithm.
5. Assigning non-zero standard formation enthalpy to an element in its standard state.
6. Reversing a Hess's law equation without reversing the sign of its enthalpy change.
7. Confusing bond breaking and bond formation. Breaking requires energy; formation releases it.
8. Assuming an exothermic reaction must be spontaneous. Entropy and temperature also matter.
9. Using irreversible heat in the reversible entropy relation.
10. Confusing spontaneity with reaction rate. Thermodynamics predicts feasibility, not speed.

Quick Revision Notes for Chemical Thermodynamics

• Open systems exchange energy and matter; closed systems exchange energy only; isolated systems exchange neither.
• Internal energy, enthalpy and entropy are state functions.
• Heat and work are path functions.
• The first law is conservation of energy for thermodynamic systems.
• At constant volume, pressure-volume work is zero.
• At constant pressure, heat exchanged equals enthalpy change.
• Exothermic reactions have negative enthalpy change, while endothermic reactions have positive enthalpy change.
• Standard reaction enthalpy can be calculated from standard formation enthalpies.
• Hess's law allows reactions to be added, reversed and scaled.
• Kirchhoff's equation accounts for the effect of temperature on reaction enthalpy.
• Entropy generally increases from solid to liquid to gas.
• Gibbs free energy determines thermodynamic spontaneity at constant temperature and pressure.
• Standard Gibbs free energy connects thermodynamics with the equilibrium constant.

Problem-solving routine: Start Chemical Thermodynamics questions by identifying the thermodynamic quantity being asked and the conditions of the process. Check the sign convention carefully, keep energy units consistent and determine whether the problem is based on the first law, enthalpy, Hess's law, entropy or Gibbs free energy before substituting values.