Revision Notes: Thermodynamics | Chemistry Class 11 - NEET PDF Download

Introduction

Thermodynamics studies energy transformations in chemical and physical processes, focusing on macroscopic systems in equilibrium. It quantifies energy changes as heat (q) and work (w), addressing questions like: How do we determine energy changes in reactions? What drives them? To what extent do they proceed?

Thermodynamic Terms

• System and Surroundings: System is the part of the universe under study; surroundings are everything else interacting with it. Universe = System + Surroundings.
• Types of Systems:
  • Open: Exchanges energy and matter (e.g., reactants in an open beaker).
  • Closed: Exchanges energy, not matter (e.g., reactants in a sealed copper vessel).
  • Isolated: No exchange of energy or matter (e.g., reactants in a thermos).
• State of the System: Defined by state functions (e.g., pressure (p), volume (V), temperature (T), amount (n)), which depend only on the system’s state, not its path.
• Internal Energy (U): Total energy of the system (chemical, mechanical, etc.), a state function. Changes when heat is transferred, work is done, or matter enters/leaves.

First Law of Thermodynamics

• Statement: Energy of an isolated system is constant; energy can neither be created nor destroyed.
• Mathematical Form: ΔU = q + w, where ΔU is the change in internal energy, q is heat (positive if absorbed), and w is work (positive if done on the system).
• Work (w):
  • Pressure-volume work: w = -p_ex ΔV (negative for expansion, positive for compression).
  • Reversible isothermal: w_rev = -nRT ln(V_f / V_i).
  • Free expansion: w = 0 (p_ex = 0).Work done on an ideal gas in acylinder when it is compressed bya constant external pressure, p_ex(in single step) is equal to the shadedarea.
• Heat (q): Energy transfer due to temperature difference; q = C ΔT (C is heat capacity).

Enthalpy (H)

• Definition: H = U + pV, a state function suitable for constant pressure processes.
• Change: ΔH = ΔU + pΔV = q_p (heat at constant pressure). ΔH = ΔU + Δn_g RT for gases.
• Sign: Negative for exothermic (heat evolved), positive for endothermic (heat absorbed).

Properties

• Extensive: Depend on amount (e.g., U, H, V, mass).
• Intensive: Independent of amount (e.g., T, p, density).
• Heat Capacity: C = q / ΔT; molar (C_m = C/n), specific (c = C/m). For ideal gas: C_p - C_v = R.

Calorimetry

• ΔU Measurement: Bomb calorimeter (constant volume); q_v = ΔU, no work (ΔV = 0).
• ΔH Measurement: Constant pressure calorimeter; q_p = ΔH.

 Enthalpy Change of Reactions (Δ_rH)

• Definition: Δ_rH = Σ a_i H_products - Σ b_i H_reactants.
• Standard State: Pure form at 1 bar, 298 K (ΔH^°).
• Types:
  • Fusion (Δ_fusH^°): Melting 1 mol solid (e.g., H₂O(s) → H₂O(l), 6.01 kJ/mol).
  • Vaporization (Δ_vapH^°): Vaporizing 1 mol liquid (e.g., H₂O(l) → H₂O(g), 40.79 kJ/mol).
  • Sublimation (Δ_subH^°): Subliming 1 mol solid (e.g., CO₂(s), 25.2 kJ/mol).
  • Formation (Δ_fH^°): Forming 1 mol from elements in standard states (e.g., CH₄, -74.81 kJ/mol).
• Standard Enthalpy Changes:

Hess’s Law

• Statement: Enthalpy change of a reaction is the same regardless of the path, if initial and final states are identical.
• Application: Calculate ΔH using formation enthalpies or intermediate steps.

Spontaneity

• Spontaneous Process: Occurs naturally (e.g., exothermic reactions).
• Non-spontaneous: Requires external energy.
• Entropy (S): Measure of disorder; ΔS_total > 0 for spontaneity.
• Gibbs Energy (G): ΔG = ΔH - TΔS; ΔG < 0 (spontaneous), ΔG > 0 (non-spontaneous), ΔG = 0 (equilibrium).

Laws of Thermodynamics

• Second Law: Entropy of an isolated system increases in a spontaneous process.
• Third Law: Entropy of a pure crystalline substance is zero at 0 K.

Gibbs Energy and Equilibrium

• Relation: Δ_rG^° = -RT ln K; K > 1 (exothermic, spontaneous), K < 1 (endothermic, less spontaneous).
• Effect of Temperature:

Summary

Thermodynamics quantifies energy changes via the first law (ΔU = q + w), introducing state functions U and H. Work (pressure-volume) and heat are path-dependent, but ΔU and ΔH depend only on initial and final states. Calorimetry measures ΔU (bomb) and ΔH (constant pressure). Reaction enthalpies (Δ_rH) include formation, fusion, etc., calculated using Hess’s law. Spontaneity depends on entropy (S) and Gibbs energy (ΔG), with ΔG = 0 at equilibrium, linked to the equilibrium constant (K).