CBSE Class 11 Chemistry Chapter 6: Thermodynamics Notes

Introduction to Thermodynamic Systems and Properties

Thermodynamics is the study of energy transformations and its relation to macroscopic properties of matter. It helps predict the feasibility and direction of a process.

System, Surroundings, and Boundary

A system is the part of the universe under thermodynamic investigation. The surroundings constitute everything else in the universe outside the system. A boundary separates the system from its surroundings. Boundaries can be real or imaginary, rigid or flexible, and adiabatic (no heat exchange) or diathermic (allows heat exchange).

Types of Systems

• Open System: Exchanges both energy (heat and work) and matter with its surroundings (e.g., an open cup of hot coffee).
• Closed System: Exchanges energy but not matter with its surroundings (e.g., a sealed flask of hot coffee).
• Isolated System: Exchanges neither energy nor matter with its surroundings (e.g., a perfectly insulated thermos flask).

State Functions vs. Path Functions

• State Functions: Properties whose values depend only on the initial and final states of the system, irrespective of the path taken (e.g., pressure (P), volume (V), temperature (T), internal energy (U), enthalpy (H), entropy (S), Gibbs free energy (G)). They are exact differentials.
• Path Functions: Properties whose values depend on the path taken to change the state of the system (e.g., heat (q), work (w)). They are inexact differentials.

Intensive vs. Extensive Properties

• Intensive Properties: Independent of the amount of matter present in the system (e.g., temperature, pressure, density, refractive index, specific heat, molar heat capacity).
• Extensive Properties: Depend on the amount of matter present in the system (e.g., mass, volume, internal energy, enthalpy, entropy, Gibbs free energy, heat capacity).

Key Thermodynamic Terms

Internal Energy (U)

The total energy contained within a thermodynamic system, including kinetic and potential energies of its constituent particles. It is a state function.

Enthalpy (H)

A thermodynamic potential that is the sum of the internal energy (U) and the product of pressure (P) and volume (V) of a system. H = U + PV. It represents the heat absorbed or released at constant pressure.

Entropy (S)

A measure of the randomness or disorder of a system. The greater the disorder, the higher the entropy. It is a state function.

Gibbs Free Energy (G)

A thermodynamic potential that measures the 'useful' or process-initiating work obtainable from an isothermal, isobaric thermodynamic system. G = H - TS. It determines the spontaneity of a process.

Spontaneous Process

A process that can occur without any external input of energy once initiated, under a given set of conditions. It proceeds in a definite direction.

Heat Capacity (C)

The amount of heat required to raise the temperature of a substance by one degree Celsius (or Kelvin). It can be at constant volume (Cv) or constant pressure (Cp).

Exothermic Reaction

A chemical reaction that releases heat to its surroundings, resulting in a negative change in enthalpy (∆H < 0).

Endothermic Reaction

A chemical reaction that absorbs heat from its surroundings, resulting in a positive change in enthalpy (∆H > 0).

Must Remember: Fundamental Laws and Formulas

• First Law of Thermodynamics (Law of Conservation of Energy): ∆U = q + w, where ∆U is change in internal energy, q is heat, and w is work. (Sign convention: q positive if heat absorbed by system, w positive if work done on system).
• Work done (w): For an irreversible expansion, w = -P_ext∆V. For a reversible isothermal expansion of an ideal gas, w = -nRT ln(V₂/V₁).
• Enthalpy (H): H = U + PV. For a reaction at constant pressure, ∆H = q_p. The relationship between ∆H and ∆U is ∆H = ∆U + ∆n_gRT, where ∆n_g is (moles of gaseous products - moles of gaseous reactants).
• Hess's Law of Constant Heat Summation: The total enthalpy change for a reaction is independent of the path by which the reaction occurs, as long as the initial and final conditions are the same.
• Second Law of Thermodynamics: For a spontaneous process in an isolated system, the entropy of the universe always increases (∆S_universe > 0). For a spontaneous process, ∆S_total = ∆S_system + ∆S_surroundings > 0.
• Gibbs Free Energy (G): G = H - TS. The change in Gibbs free energy (∆G) at constant temperature and pressure determines spontaneity: ∆G < 0: Process is spontaneous. ∆G = 0: Process is at equilibrium. * ∆G > 0: Process is non-spontaneous (reverse process is spontaneous).
• Third Law of Thermodynamics: The entropy of a perfectly crystalline substance at absolute zero (0 K) is taken as zero.
• Relationship between ∆G° and K_eq: ∆G° = -RT ln K_eq, where K_eq is the equilibrium constant. A large positive K_eq implies a negative ∆G° and spontaneity.

Reversible vs. Irreversible Processes

Worked Examples

• {"title":"1. Calculating ∆H from ∆U","bodyMarkdown":"Problem: For the reaction N₂(g) + 3H₂(g) → 2NH₃(g) at 298 K, ∆U = -92.38 kJ/mol. Calculate ∆H.\n\nSolution:\n1. Identify moles of gaseous products (n_products) = 2 moles NH₃(g).\n2. Identify moles of gaseous reactants (n_reactants) = 1 mole N₂(g) + 3 moles H₂(g) = 4 moles.\n3. Calculate ∆n_g = n_products - n_reactants = 2 - 4 = -2 mol.\n4. Use the relation: ∆H = ∆U + ∆n_gRT\n Given R = 8.314 J/K mol = 0.008314 kJ/K mol, T = 298 K.\n ∆H = -92.38 kJ + (-2 mol)(0.008314 kJ/K mol)(298 K)\n ∆H = -92.38 kJ - 4.95 kJ\n ∆H = -97.33 kJ/mol"}

Exam Traps & Key Focus Areas

1. Sign Conventions: Always pay close attention to the sign conventions for heat (q) and work (w). Heat absorbed by the system is positive (+q), heat released is negative (-q). Work done on the system is positive (+w), work done by the system is negative (-w). Mistakes here are very common.
2. Units: Ensure consistent units, especially when using the gas constant R (J/K mol or kJ/K mol). Convert all energy terms (∆U, ∆H, ∆G) to the same unit (kJ or J).
3. Spontaneity Criteria: Memorize the conditions for spontaneity based on ∆G, ∆H, and ∆S (∆G = ∆H - T∆S). Understand how temperature affects spontaneity, especially when ∆H and ∆S have the same sign.
4. Hess's Law: Practice applying Hess's Law accurately. Remember that if you reverse a reaction, the sign of ∆H changes. If you multiply a reaction by a factor, multiply ∆H by the same factor.
5. State vs. Path Functions: Clearly distinguish between these. Internal energy, enthalpy, entropy, and Gibbs free energy are state functions; heat and work are path functions.

Practice Questions with Solutions

• Q: What is the main difference between an open system and a closed system? A: An open system exchanges both matter and energy with its surroundings, while a closed system exchanges only energy but not matter.
• Q: Under what conditions is ∆H equal to ∆U? A: ∆H = ∆U when there is no change in the number of moles of gaseous substances in a reaction (∆n_g = 0), or when the reaction occurs in a closed container at constant volume (∆V = 0).
• Q: What does a negative value of Gibbs free energy change (∆G) signify for a process? A: A negative ∆G indicates that the process is spontaneous under the given constant temperature and pressure conditions.
• Q: How is the entropy of the universe affected during a spontaneous process? A: According to the Second Law of Thermodynamics, the entropy of the universe (∆S_universe) increases during any spontaneous process (∆S_universe > 0).

Frequently Asked Questions