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Class 11 Chemistry: Thermodynamics One-Shot Notes

Comprehensive revision notes covering basic thermodynamic terms, systems, state variables, extensive and intensive properties, and thermodynamic processes. Explains work done, heat capacity, enthalpy relations, Hess's law, the Born-Haber cycle, entropy, and Gibbs free energy for spontaneity.

Class 11 Science
JEE
NEET
Chemistry
English · 11 topics
Checked against the video on 25 Oct 2026

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Class 11 Chemistry: Thermodynamics One-Shot Notes

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In short

Comprehensive revision notes covering basic thermodynamic terms, systems, state variables, extensive and intensive properties, and thermodynamic processes. Explains work done, heat capacity, enthalpy relations, Hess's law, the Born-Haber cycle, entropy, and Gibbs free energy for spontaneity.

Introduction to Thermodynamics and Systems

Thermodynamics deals with the study of relations between heat, work, temperature, and energy transformations into other forms.
Macroscopic quantities such as volume, pressure, temperature, internal energy, entropy, and enthalpy are studied in thermodynamics.
Must knowThe system refers to that part of the universe in which observations are made, while the remaining universe constitutes the surroundings.
The universe is defined by the relation: Universe = System + Surroundings.
The boundary that separates the system from its surroundings is called the system boundary.

Types of Thermodynamic Systems

Must knowAn open system is one in which both matter and energy can be exchanged between the system and its surroundings.
A closed system allows the transfer of energy across its boundary, but no transfer of matter takes place.
An isolated system permits neither the transfer of matter nor the transfer of energy between the system and the surroundings (e.g., an ideal thermos flask).

System Type	Mass Transfer	Energy Transfer
Open	Yes	Yes
Closed	No	Yes
Isolated	No	No

State Variables, Extensive and Intensive Properties

State variables are parameters that describe the equilibrium state of a system (such as pressure, volume, temperature, mass, and composition).
The change in a state function depends only on the initial and final conditions of the system, not on the path taken.
Must knowAn extensive property depends upon the total amount of matter present in the system (e.g., mass, volume, moles, internal energy).
Must knowAn intensive property is independent of the size or amount of matter in the system (e.g., pressure, temperature, concentration/molarity, density).
The ratio of two extensive properties results in an intensive property (e.g., Molarity = moles / volume).

Thermodynamic Processes

An isothermal process occurs at a constant temperature (T = constant), where PV = constant.
An isobaric process occurs at constant pressure (P is fixed), where volume and temperature change with time.
An isochoric process occurs at constant volume (V is constant); therefore, no work is done on or by the gas (W = 0).
Must knowAn adiabatic process is one in which there is no change in heat (heat exchange is zero, ΔQ = 0 or q = 0); the system is enclosed by thermally insulated walls.
In any cyclic process, the change in a state function over the complete cycle is zero.
Heat (q) and work (w) are path functions, whereas internal energy (U) is a state function.

Process	Constant Parameter	Condition
Isothermal	Temperature	ΔT = 0
Isobaric	Pressure	ΔP = 0
Isochoric	Volume	ΔV = 0, W = 0
Adiabatic	Heat	q = 0

Internal Energy, Work, and First Law of Thermodynamics

Internal energy (U) is the total energy stored within a substance, consisting of various forms of molecular energy, and is a state function.
If internal energy decreases (energy released), ΔU is negative; if energy is absorbed, ΔU is positive.
Under adiabatic conditions where heat cannot exchange, the change in internal energy equals the adiabatic work: ΔU = W_ad.
Sign convention: Heat absorbed by the system is positive (+q); heat given out is negative (-q). Work done on the system is positive (+w); work done by the system is negative (-w).
Must knowThe First Law of Thermodynamics states that the energy of an isolated system is constant, expressed mathematically as: ΔU = q + w.

Pressure-Volume Work and Reversible Expansion

Work done during gas expansion or compression under external pressure is given by: W = -P_ext × ΔV.
For variable pressure, work is calculated by integration: W = - ∫(from V1 to V2) P dV.
On a P-V graph, work done equals the area under the P-V curve.
Must knowWhen a gas expands (V2 > V1), ΔV is positive, so work done by the gas is negative (W is negative). When a gas is compressed (V2 < V1), ΔV is negative, so work done on the gas is positive (W is positive).
Reversible processes proceed infinitely slowly through infinitesimal steps maintaining near-equilibrium, whereas irreversible processes occur spontaneously in finite time.
Must knowThe work done in a reversible isothermal expansion of an ideal gas is: W_rev = -2.303 × n × R × T × log10(V_f / V_i).
A reversible expansion provides more work output (larger area under the curve) than an irreversible expansion between the same limits.

Heat Capacity and Relation between Cp and Cv

Heat capacity (C) is the amount of heat required to raise the temperature of a system by 1 unit: C = ΔQ / ΔT (or Q = C × ΔT); its unit is J/K.
Specific heat capacity (c) is the heat required to raise the temperature of one unit mass of a substance by 1°C: Q = m × c × ΔT; unit is J/(kg·K).
Molar heat capacity is the heat required to raise the temperature of 1 mole of a substance by 1 unit: Q = n × C × ΔT; unit is J/(K·mol).
Heat capacity at constant volume is denoted as C_v (where q_v = ΔU), and at constant pressure as C_p (where q_p = ΔH).
Must knowThe relation between molar heat capacity at constant pressure and constant volume for an ideal gas is: C_p = C_v + R (or C_p - C_v = R).

Enthalpy and Enthalpy Changes

Enthalpy (H) is the total energy content of a system combining internal energy and pressure-volume energy: H = U + PV, or ΔH = ΔU + PΔV at constant pressure.
At constant pressure, heat exchanged equals enthalpy change: q_p = ΔH.
Must knowFor an exothermic reaction, heat is evolved and ΔH is negative; for an endothermic reaction, heat is absorbed and ΔH is positive.
For ideal gases, enthalpy change relates to internal energy change via: ΔH = ΔU + Δn_g × R × T.
Enthalpy of atomization is the enthalpy change upon breaking one mole of a substance completely into atoms in the gaseous state.
Lattice enthalpy is the energy required to completely separate one mole of a solid ionic compound into gaseous constituent ions.

Hess's Law and Born-Haber Cycle

Must knowHess's Law of Constant Heat Summation states that if a reaction takes place in several steps, the overall enthalpy change is the sum of the enthalpy changes of the individual steps.
The Born-Haber cycle is an indirect cyclic method based on Hess's Law used to calculate lattice energy from experimental data (sublimation, ionization, dissociation, electron affinity, and formation enthalpies).
In the formation of NaCl(s), ΔH_f° equals the sum of ΔH_sub (Na), Ionization Energy (Na), 1/2 ΔH_diss (Cl2), Electron Affinity (Cl), and Lattice Energy (U).

Entropy, Second and Third Laws of Thermodynamics

A spontaneous process occurs on its own without requiring an external driving force and possesses a natural tendency to proceed in one direction.
Must knowEntropy (S) is a measure of the randomness or disorder of a system; it is a state function.
The order of entropy among physical states is: Solid (minimum entropy) < Liquid < Gas (maximum entropy).
Entropy change relates to heat and temperature by: ΔS = q_rev / T.
Must knowFor a reversible process, ΔS_total = 0; for an irreversible (spontaneous) process, ΔS_total > 0 (Second Law of Thermodynamics).
Must knowThe Third Law of Thermodynamics states that the entropy of a perfectly crystalline substance approaches zero as the absolute temperature approaches zero (0 Kelvin).

Gibbs Free Energy and Spontaneity

Gibbs free energy (G) is a state function representing the maximum amount of useful work obtainable from a system at constant temperature and pressure: G = H - TS.
Must knowThe Gibbs-Helmholtz equation defines the change in free energy: ΔG = ΔH - TΔS.
Must knowCriteria for spontaneity: ΔG < 0 indicates a spontaneous process; ΔG > 0 indicates a non-spontaneous process; ΔG = 0 indicates the system is at equilibrium.
The standard free energy change relates to the equilibrium constant (K) by: ΔG° = -RT ln K = -2.303 RT log10 K.
When ΔG° < 0, K > 1 (products favored); when ΔG° > 0, K < 1 (reactants favored).

ΔH	ΔS	ΔG	Process Nature
Negative	Positive	Always Negative	Spontaneous at all temperatures
Positive	Negative	Always Positive	Non-spontaneous at all temperatures
Positive	Positive	Negative at high T	Spontaneous only at high temperatures
Negative	Negative	Negative at low T	Spontaneous only at low temperatures

Key terms

System: The specific part of the universe selected for thermodynamic study and observation.
Surroundings: Everything in the universe outside the system.
State Function: A thermodynamic property whose value depends only on the current state of the system, not on how that state was reached.
Extensive Property: A property that depends on the quantity or size of matter present in the system.
Intensive Property: A property that is independent of the amount or mass of matter in the system.
Adiabatic Process: A process conducted without any heat transfer between the system and its surroundings (q = 0).
Isochoric Process: A process occurring at constant volume where work done is zero.
Isobaric Process: A process occurring at constant pressure.
Enthalpy (H): The heat content of a system at constant pressure, defined as U + PV.
Entropy (S): A thermodynamic measure of molecular disorder or randomness within a system.
Gibbs Free Energy (G): Thermodynamic capacity of a system to perform useful work at constant T and P.
Hess's Law: Principle stating that the total enthalpy change for a chemical reaction is independent of the pathway taken.

Formulas

ΔU = q + w: First Law of Thermodynamics
W = -P_ext × ΔV: Work done at constant external pressure
W_rev = -2.303 × n × R × T × log10(V_f / V_i): Reversible isothermal work of an ideal gas
C_p - C_v = R: Mayer's relation for ideal gas molar heat capacities
ΔH = ΔU + PΔV: Enthalpy change relation at constant pressure
ΔS = q_rev / T: Entropy change formula
ΔG = ΔH - TΔS: Gibbs-Helmholtz equation for spontaneity
ΔG° = -2.303 × R × T × log10 K: Standard free energy and equilibrium constant relation

Exam tips and traps

Remember that work done BY the system (expansion) is negative in chemistry (W = -PΔV), which is opposite to the usual convention in physics.
A reaction is always spontaneous at all temperatures when ΔH is negative (exothermic) and ΔS is positive (increasing disorder).

Quick revision

Read these the night before the exam.

An isolated system cannot exchange either energy or matter with its surroundings.
Temperature, pressure, and density are intensive properties; mass and volume are extensive properties.
Work done in an isochoric process is always zero because ΔV = 0.
Internal energy and enthalpy are state functions, whereas heat and work are path functions.
For an isothermal expansion of an ideal gas, the change in internal energy (ΔU) is zero.
The relation between heat capacities is C_p = C_v + R for one mole of an ideal gas.
An exothermic reaction has a negative ΔH, while an endothermic reaction has a positive ΔH.
Hess's Law states that total enthalpy change equals the sum of enthalpy changes of individual steps.
The third law of thermodynamics states that entropy of a pure crystalline substance is zero at 0 K.
A negative value of ΔG is the necessary condition for a spontaneous process at constant T and P.

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