Class 11 Chemistry Notes Chapter 6 (Chapter 6) – Examplar Problems (English) Book

Detailed Notes with MCQs of Chapter 6: Thermodynamics from your NCERT Class 11 Chemistry Exemplar. This is a crucial chapter, not just for your Class 11 exams, but also forms the foundation for many concepts tested in various government exams. Pay close attention to the definitions, laws, and formulas.

Chapter 6: Thermodynamics - Detailed Notes for Competitive Exams

1. Introduction & Basic Terminology

• System: The specific part of the universe under investigation (e.g., reactants in a beaker).
• Surroundings: Everything else in the universe outside the system.
• Boundary: The real or imaginary surface separating the system from the surroundings.
• Types of Systems:
  • Open System: Exchanges both energy and matter with the surroundings (e.g., boiling water in an open beaker).
  • Closed System: Exchanges energy but not matter with the surroundings (e.g., boiling water in a sealed container).
  • Isolated System: Exchanges neither energy nor matter with the surroundings (e.g., thermos flask - an approximation).

2. Thermodynamic Properties

• State Functions (State Variables): Properties whose value depends only on the current state of the system, independent of the path taken to reach that state. Examples: Pressure (P), Volume (V), Temperature (T), Internal Energy (U), Enthalpy (H), Entropy (S), Gibbs Free Energy (G). Change in state function (ΔX) depends only on initial and final states (ΔX = X_final - X_initial).
• Path Functions: Properties whose value depends on the path followed during a process. Examples: Heat (q), Work (w).
• Intensive Properties: Properties independent of the amount of substance present. Examples: Temperature (T), Pressure (P), Density (ρ), Molar Heat Capacity, Concentration.
• Extensive Properties: Properties dependent on the amount of substance present. Examples: Mass (m), Volume (V), Internal Energy (U), Enthalpy (H), Entropy (S), Gibbs Free Energy (G), Heat Capacity.

3. Thermodynamic Processes

• Isothermal Process: Temperature remains constant (ΔT = 0). For ideal gas, ΔU = 0.
• Adiabatic Process: No heat exchange between system and surroundings (q = 0).
• Isobaric Process: Pressure remains constant (ΔP = 0).
• Isochoric Process: Volume remains constant (ΔV = 0). Work done (w = -P_ext ΔV) is zero.
• Reversible Process: Process carried out infinitesimally slowly such that the system remains in equilibrium with the surroundings at each stage. Can be reversed by an infinitesimal change in conditions. Maximum work is obtained.
• Irreversible Process: Process carried out rapidly; system is not in equilibrium with surroundings. Cannot be reversed exactly. Work done is less than in a reversible process. All natural processes are irreversible.

4. Internal Energy (U)

• Sum of all forms of energy associated with a system (kinetic, potential, electronic, nuclear, etc.).
• It's an extensive property and a state function.
• Absolute value cannot be determined, but change (ΔU) can be measured.
• For an ideal gas, U depends only on temperature.

5. First Law of Thermodynamics (Law of Conservation of Energy)

• Energy can neither be created nor destroyed, only converted from one form to another.
• Mathematical Statement: ΔU = q + w
  • ΔU = Change in internal energy
  • q = Heat absorbed by the system
  • w = Work done on the system
• Sign Conventions (Important!):
  • Heat absorbed by the system: q = +ve
  • Heat released by the system: q = -ve
  • Work done on the system: w = +ve
  • Work done by the system: w = -ve

6. Work (w)

• Pressure-Volume Work: Work done during expansion or compression against external pressure.
  • w = -P_ext ΔV = -P_ext (V_f - V_i)
  • Expansion: V_f > V_i, ΔV = +ve, w = -ve (Work done by the system)
  • Compression: V_f < V_i, ΔV = -ve, w = +ve (Work done on the system)
• Work in Free Expansion (Expansion into vacuum): P_ext = 0, so w = 0.
• Work in Isothermal Reversible Expansion (Ideal Gas):
  • w_rev = -nRT ln(V_f / V_i) = -2.303 nRT log(V_f / V_i)
  • Since P₁V₁ = P₂V₂, w_rev = -2.303 nRT log(P_i / P_f)
• Work in Isochoric Process: ΔV = 0, so w = 0.

7. Heat (q) and Heat Capacity

• Heat Capacity (C): Heat required to raise the temperature of a substance by 1°C or 1K. C = q / ΔT. (Extensive property)
• Molar Heat Capacity (C_m): Heat required to raise the temperature of 1 mole of a substance by 1°C or 1K. C_m = C / n = q / (n ΔT). (Intensive property)
• Specific Heat Capacity (c): Heat required to raise the temperature of 1 gram of a substance by 1°C or 1K. c = C / m = q / (m ΔT). (Intensive property)
• Heat Capacity at Constant Volume (C_v): C_v = (q_v) / ΔT. Since ΔV=0, w=0. From 1st Law, ΔU = q_v. So, C_v = (ΔU / ΔT)_v.
• Heat Capacity at Constant Pressure (C_p): C_p = (q_p) / ΔT.
• Relationship for Ideal Gas: C_p - C_v = R (R = Gas constant)
• Ratio of Heat Capacities (γ): γ = C_p / C_v

8. Enthalpy (H)

• Total heat content of the system at constant pressure.
• Definition: H = U + PV
• It's an extensive property and a state function.
• Change in Enthalpy (ΔH): ΔH = ΔU + Δ(PV)
• At constant pressure: ΔH = ΔU + PΔV
• Since ΔU = q + w = q - P_extΔV. If P_ext = P (constant pressure), then ΔU = q_p - PΔV, or q_p = ΔU + PΔV.
• Therefore, ΔH = q_p (Heat change at constant pressure).
• Relationship between ΔH and ΔU for Gaseous Reactions:
  • ΔH = ΔU + Δn_g RT
  • Δn_g = (Total moles of gaseous products) - (Total moles of gaseous reactants)
  • If Δn_g = 0, then ΔH = ΔU.
  • If Δn_g > 0, then ΔH > ΔU.
  • If Δn_g < 0, then ΔH < ΔU.

9. Enthalpy Changes (Thermochemistry)

• Standard State: 1 bar pressure (or 1 atm historically) and specified temperature (usually 298 K). Represented by superscript '°'.
• Standard Enthalpy of Reaction (Δ_r H°): Enthalpy change when reaction occurs between substances in their standard states.
• Standard Enthalpy of Formation (Δ_f H°): Enthalpy change when 1 mole of a compound is formed from its constituent elements in their most stable standard states (reference states). Δ_f H° of elements in their reference state is zero (e.g., O₂(g), C(graphite), H₂(g)).
  • Δ_r H° = Σ ν_p Δ_f H°(Products) - Σ ν_r Δ_f H°(Reactants) (ν = stoichiometric coefficients)
• Standard Enthalpy of Combustion (Δ_c H°): Enthalpy change when 1 mole of a substance undergoes complete combustion in excess oxygen at standard conditions. Usually exothermic (Δ_c H° is negative).
• Enthalpy of Phase Transition:
  • Enthalpy of Fusion (Δ_fus H°): Enthalpy change when 1 mole of solid melts into liquid at its melting point and standard pressure.
  • Enthalpy of Vaporization (Δ_vap H°): Enthalpy change when 1 mole of liquid vaporizes into gas at its boiling point and standard pressure.
  • Enthalpy of Sublimation (Δ_sub H°): Enthalpy change when 1 mole of solid sublimes into gas at a given temperature below its melting point. Δ_sub H° = Δ_fus H° + Δ_vap H°
• Hess's Law of Constant Heat Summation: The total enthalpy change for a reaction is the same, whether the reaction takes place in one step or in several steps. Allows calculation of enthalpy changes for reactions that cannot be measured directly.
• Bond Dissociation Enthalpy (Bond Energy): Energy required to break 1 mole of a specific type of bond in gaseous molecules. Always positive (endothermic).
  • For diatomic molecules, it's simply the enthalpy change for dissociation.
  • For polyatomic molecules, average bond enthalpy is used.
  • Δ_r H° ≈ Σ (Bond Enthalpies)_Reactants - Σ (Bond Enthalpies)_Products (Note: Reactants minus Products for bond enthalpies)
• Lattice Enthalpy (Δ_lattice H°): Enthalpy change when 1 mole of an ionic compound dissociates into its gaseous ions. Can be calculated using the Born-Haber cycle (application of Hess's Law).

10. Spontaneity

• Spontaneous Process: A process that has a natural tendency to occur under given conditions without external intervention. May be fast or slow. Irreversible.
• Driving Forces for Spontaneity:
  • Tendency to achieve minimum energy (negative ΔH favors spontaneity).
  • Tendency to achieve maximum randomness (positive ΔS favors spontaneity).

11. Entropy (S)

• Measure of the degree of randomness or disorder of a system.
• It's an extensive property and a state function.
• Units: J K⁻¹ mol⁻¹.
• Entropy increases with: Solid < Liquid < Gas; increase in temperature; increase in number of moles (especially gas); mixing of substances.
• Change in Entropy: ΔS = q_rev / T (at constant T)
• Second Law of Thermodynamics: The entropy of the universe (system + surroundings) always increases during a spontaneous (irreversible) process.
  • ΔS_total = ΔS_universe = ΔS_system + ΔS_surroundings > 0 (for spontaneous process)
  • ΔS_total = 0 (for reversible process/equilibrium)
• For surroundings: ΔS_surr = -ΔH_sys / T (at constant P, T)

12. Gibbs Free Energy (G)

• A thermodynamic potential that combines enthalpy and entropy to determine spontaneity under constant temperature and pressure.

• Definition: G = H - TS

• It's an extensive property and a state function.

• Change in Gibbs Energy: ΔG = ΔH - TΔS (Gibbs-Helmholtz Equation, at constant T)

• Criterion for Spontaneity (at constant T and P):

  • ΔG < 0: Process is spontaneous.
  • ΔG > 0: Process is non-spontaneous (reverse process is spontaneous).
  • ΔG = 0: System is in equilibrium.

• Effect of Temperature on Spontaneity: The sign of ΔG depends on the signs of ΔH and ΔS, and the temperature T.

  ΔH	ΔS	ΔG = ΔH - TΔS	Spontaneity
  -ve	+ve	Always -ve	Spontaneous at all temperatures
  +ve	-ve	Always +ve	Non-spontaneous at all temperatures
  -ve	-ve	-ve at low T, +ve at high T	Spontaneous at low T, Non-spontaneous at high T
  +ve	+ve	+ve at low T, -ve at high T	Non-spontaneous at low T, Spontaneous at high T

• Standard Gibbs Energy Change (ΔG°): Gibbs energy change when reactants in standard states convert to products in standard states.

• Relationship between ΔG° and Equilibrium Constant (K):

  • ΔG = ΔG° + RT ln Q (Q = Reaction Quotient)
  • At equilibrium, ΔG = 0 and Q = K.
  • 0 = ΔG° + RT ln K
  • ΔG° = -RT ln K = -2.303 RT log K
  • If K > 1, ΔG° < 0 (Products favored, spontaneous under standard conditions).
  • If K < 1, ΔG° > 0 (Reactants favored, non-spontaneous under standard conditions).
  • If K = 1, ΔG° = 0 (Equilibrium under standard conditions).

13. Third Law of Thermodynamics

• The entropy of a perfectly crystalline solid approaches zero as the temperature approaches absolute zero (0 K).
• Allows calculation of absolute entropies of substances at different temperatures.

Multiple Choice Questions (MCQs)

1. Which of the following is NOT a state function?
   a) Internal Energy (U)
   b) Enthalpy (H)
   c) Work (w)
   d) Entropy (S)

2. For an adiabatic process, which condition is correct?
   a) ΔT = 0
   b) ΔP = 0
   c) q = 0
   d) w = 0

3. The first law of thermodynamics is essentially the law of:
   a) Conservation of Mass
   b) Conservation of Energy
   c) Conservation of Momentum
   d) Constant Entropy

4. For the reaction N₂(g) + 3H₂(g) ⇌ 2NH₃(g), the relationship between ΔH and ΔU is:
   a) ΔH = ΔU
   b) ΔH = ΔU - 2RT
   c) ΔH = ΔU + 2RT
   d) ΔH = ΔU - RT

5. According to Hess's Law, the enthalpy change of a reaction depends only on:
   a) The path taken
   b) The initial and final states
   c) The temperature of the reaction
   d) The pressure of the reaction

6. For a process to be spontaneous at constant temperature and pressure, the condition is:
   a) ΔS > 0
   b) ΔH < 0
   c) ΔG = 0
   d) ΔG < 0

7. The standard enthalpy of formation (Δ_f H°) of an element in its most stable reference state is:
   a) Always positive
   b) Always negative
   c) Zero
   d) Depends on the element

8. In the isothermal reversible expansion of an ideal gas, which statement is TRUE?
   a) ΔU = 0 and w = 0
   b) ΔU > 0 and q = -w
   c) ΔU = 0 and q = -w
   d) ΔU < 0 and q = w

9. If ΔH is positive and ΔS is positive for a reaction, the reaction will be spontaneous:
   a) At high temperatures
   b) At low temperatures
   c) At all temperatures
   d) Never spontaneous

10. The relationship between standard free energy change (ΔG°) and equilibrium constant (K) is:
    a) ΔG° = RT ln K
    b) ΔG° = -RT ln K
    c) ΔG° = RT log K
    d) K = -RT ln ΔG°

Answer Key for MCQs:

1. c) Work (w)
2. c) q = 0
3. b) Conservation of Energy
4. b) ΔH = ΔU - 2RT (Δn_g = 2 - (1+3) = -2)
5. b) The initial and final states
6. d) ΔG < 0
7. c) Zero
8. c) ΔU = 0 and q = -w (Since ΔU = q + w, if ΔU=0 then q = -w)
9. a) At high temperatures (The TΔS term needs to overcome ΔH)
10. b) ΔG° = -RT ln K

Make sure you understand the reasoning behind each answer. Go through these notes thoroughly, practice problems from the Exemplar book, and focus on applying these concepts. Good luck with your preparation!