Unit 4

General and Physical Chemistry

Class 12 Chemistry

Thermodynamics

Class 12 Chemistry – Thermodynamics Notes PDF

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Introduction

DefinitionThermodynamics is the study of energy, heat, work and the criteria governing energy changes and spontaneity in physical and chemical processes.

This unit covers internal energy, first law of thermodynamics, enthalpy and enthalpy changes, endothermic/exothermic processes, enthalpy of reaction/solution/formation/combustion, Laplace and Hess laws, entropy, second law, Gibbs free energy and the relationship between ΔG and equilibrium constant.

1. System, Surroundings and Energy

System

The part of the universe selected for study.

Surroundings

Everything outside the system that can interact with it.

System Type	Matter Exchange	Energy Exchange
Open	Yes	Yes
Closed	No	Yes
Isolated	No	No ideal exchange

Internal EnergyU is the total microscopic energy associated with the particles of a system. It is a state function.

Fig. 1 – System and Surroundings

2. First Law of Thermodynamics

First LawEnergy cannot be created or destroyed; it can only be transferred or converted from one form to another.

Using the chemistry sign convention where w is work done on the system:

ΔU = q + w

For pressure–volume work at constant external pressure:

w = −P_extΔV

Quantity	Positive Sign	Negative Sign
q	Heat absorbed by system	Heat released by system
w	Work done on system	Work done by system
ΔU	Internal energy increases	Internal energy decreases

ExampleA system absorbs 200 J heat and does 50 J of work on surroundings. Then q = +200 J and w = −50 J. So ΔU = 150 J.

Fig. 2 – First-Law Energy Accounting

3. Enthalpy and Enthalpy Changes

H = U + PV

At constant pressure, for processes where only pressure–volume work is relevant:

q_p = ΔH

Exothermic Process

Releases heat to surroundings; ΔH < 0.

Endothermic Process

Absorbs heat from surroundings; ΔH > 0.

Fig. 3 – Exothermic and Endothermic Energy Profiles

4. Important Enthalpy Changes

Term	Meaning
Enthalpy of reaction, ΔH_rxn	Enthalpy change when a chemical reaction occurs as specified by its stoichiometric equation.
Standard enthalpy of formation, ΔH°_f	Enthalpy change when one mole of a compound forms from its elements in their standard states.
Enthalpy of combustion, ΔH°_c	Enthalpy change when one mole of substance burns completely in oxygen under specified standard conditions.
Enthalpy of solution, ΔH_sol	Enthalpy change associated with dissolving a substance in a solvent under specified conditions.

ΔH°_rxn = ΣνΔH°_f(products) − ΣνΔH°_f(reactants)

Fig. 4 – Thermochemical Enthalpy Categories

5. Laws of Thermochemistry

Laplace Law

The heat change of a reaction has the same magnitude and opposite sign when the reaction is reversed.

A → B, ΔH = x ⇒ B → A, ΔH = −x

Hess’s Law

Hess LawThe total enthalpy change for a reaction depends only on initial and final states, not on the path taken.

Fig. 5 – Hess’s Law Path Independence

ExampleIf A → B has ΔH = +50 kJ and B → C has ΔH = −80 kJ, then A → C has ΔH = 50 + (−80) = −30 kJ.

6. Entropy and the Second Law

EntropyS is a thermodynamic state function associated with the dispersal of energy and the number of accessible microscopic arrangements of a system.

Entropy often increases when matter or energy becomes more dispersed, such as during expansion, mixing or many transitions from solid → liquid → gas.

Second Law

Second LawFor a spontaneous process, the total entropy change of the universe is positive.

ΔS_universe = ΔS_system + ΔS_surroundings > 0

Fig. 6 – Qualitative Entropy Trend

7. Gibbs Free Energy and Spontaneity

G = H − TS

At constant temperature and pressure:

ΔG = ΔH − TΔS

ΔG	Prediction
ΔG < 0	Process is thermodynamically spontaneous in the forward direction.
ΔG = 0	System is at equilibrium.
ΔG > 0	Forward process is nonspontaneous; reverse direction is favored thermodynamically.

ImportantSpontaneous does not mean fast. Thermodynamics predicts favorability, while chemical kinetics addresses reaction rate.

Fig. 7 – ΔH, ΔS and Temperature Effects on ΔG

8. Relationship Between ΔG° and Equilibrium Constant

ΔG° = −RT ln K

ΔG = ΔG° + RT ln Q

At equilibrium, ΔG = 0 and Q = K.

K	ΔG°	Interpretation
K > 1	Negative	Products favored at equilibrium.
K = 1	Zero	Neither side strongly favored under standard-state comparison.
K < 1	Positive	Reactants favored at equilibrium.

Fig. 8 – ΔG° and Equilibrium Constant

9. Worked Numerical Patterns

First LawA system absorbs 500 J heat and does 120 J work on surroundings. q = +500 J, w = −120 J, so ΔU = +380 J.

Gibbs EnergyIf ΔH = −40 kJ mol⁻¹, ΔS = −100 J mol⁻¹ K⁻¹ and T = 298 K, convert ΔS = −0.100 kJ mol⁻¹ K⁻¹. ΔG = −40 − 298(−0.100) = −40 + 29.8 = −10.2 kJ mol⁻¹.

Hess LawIf A→B has +25 kJ and B→C has −60 kJ, then A→C = +25 − 60 = −35 kJ.

EquilibriumIf K = 10 at a given T, lnK > 0, so ΔG° = −RT lnK is negative; products are favored relative to reactants at equilibrium.

10. Quick Revision & Exam Points

Important Questions

Define thermodynamics, system, surroundings and internal energy.
State and explain first law of thermodynamics.
Explain sign convention for heat and work.
Define enthalpy and distinguish endothermic and exothermic processes.
Define enthalpy of reaction, solution, formation and combustion.
State Laplace law and Hess’s law.
Define entropy and spontaneity.
State the second law of thermodynamics.
Define Gibbs free energy and use ΔG = ΔH − TΔS.
Predict spontaneity using sign of ΔG.
Explain relationship between ΔG° and equilibrium constant.

Internal energy U is a state function.
First law: ΔU = q + w.
At constant pressure, q_p = ΔH.
Exothermic: ΔH < 0.
Endothermic: ΔH > 0.
Hess law: enthalpy is path independent.
Spontaneous process has ΔS_universe > 0.
ΔG = ΔH − TΔS.
ΔG < 0 means spontaneous forward process.
At equilibrium, ΔG = 0.
ΔG° = −RT lnK.

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