Thermodynamics – Class 11 Chemistry Notes
Thermodynamics
CBSE | ISC | WBCHSE Class 11 Chemistry – Complete Notes, Principles, Laws, MCQs with Answers
1. Introduction & Basic Concepts
Thermodynamics is the branch of science that deals with the quantitative relationship between heat and other forms of energy during physical and chemical changes.
- System: Part of universe under study (e.g., a reaction vessel).
- Surroundings: Everything outside the system.
- Boundary: Separates system from surroundings—can be real or imaginary.
-
Types of Systems:
- Open: Exchange mass & energy (e.g., open beaker).
- Closed: Exchange energy only (e.g., stoppered flask).
- Isolated: No exchange (e.g., thermos flask).
- State Functions: Depend only on initial & final state (e.g., internal energy U, enthalpy H, entropy S).
- Path Functions: Depend on path/process (e.g., heat q, work w).
Knowing which quantities are state or path functions is crucial for numericals and MCQs!
2. First Law of Thermodynamics
Law of Conservation of Energy: Energy can neither be created nor destroyed; it only changes from one form to another.
- Internal energy (U): Total energy of a system (kinetic + potential).
- Change in internal energy (ΔU): ΔU = q + w
- q = Heat supplied to system (+ if absorbed, – if released)
- w = Work done on system (+) or by system (–)
ΔU = q + w
-
Work done by a gas:
- At constant pressure (expansion): w = –P_ext ΔV
- Isothermal process: ΔT = 0 (constant temperature)
- Adiabatic process: q = 0 (no heat exchange)
- Isochoric process: ΔV = 0 (constant volume); w = 0, so ΔU = q
- Isobaric process: ΔP = 0 (constant pressure)
3. Enthalpy (H), Internal Energy (U), and Heat Capacity
- Enthalpy (H): H = U + PV (useful under constant P, common in chemistry)
- Change in enthalpy (ΔH): ΔH = ΔU + Δ(PV)
- At constant pressure: ΔH = q_p (heat at constant pressure)
- At constant volume: ΔU = q_v
-
Relation (for gases):
ΔH = ΔU + ΔngRT
(Δng = change in number of gaseous moles)
- Heat capacity (C): Amount of heat required to raise T by 1 K. C = q/ΔT
- Specific heat (c): q = mcΔT
- Molar heat capacity: q = nCmΔT
4. Second & Third Laws of Thermodynamics
Second Law
- Spontaneous processes have a tendency towards greater disorder or randomness (entropy S).
- Entropy (S): Measure of disorder/randomness in a system.
- For any spontaneous process, ΔS_total = ΔS_system + ΔS_surroundings > 0
- Heat absorbed at constant T: ΔS = q_rev / T
Second Law: “All spontaneous processes increase the entropy of the universe.” — Explains the direction of heat flow and why energy conversions are never 100% efficient.
Third Law
- The entropy of a perfect crystal at absolute zero is zero: S = 0 \ (at\ 0\ K)
Important for calculations involving standard entropy values and predicting feasibility of reactions at low temperatures.
5. Thermochemical Equations & Enthalpy Changes
- Thermochemical equation: Specifies physical states and energy changes. E.g., H₂(g) + ½O₂(g) → H₂O(l); ΔH = –286 kJ/mol
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Types of Enthalpy Changes:
- ΔH_f: Enthalpy of formation
- ΔH_c: Enthalpy of combustion
- ΔH_neut: Enthalpy of neutralization
- ΔH_atom: Enthalpy of atomization
- ΔH_sub: Enthalpy of sublimation
- ΔH_vap: Enthalpy of vaporization
-
Endothermic: Heat absorbed, ΔH > 0
Exothermic: Heat released, ΔH < 0
Exothermic (ΔH < 0) and endothermic (ΔH > 0) reactions
6. Hess’s Law of Constant Heat Summation
The total enthalpy change for a reaction is the same, whether it occurs in one step or in multiple steps.
- Used to calculate enthalpy changes for reactions not possible to carry out directly.
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ΔH (overall) = ΣΔH (individual steps)
Illustration of Hess’s law
Practice: Calculate enthalpy of formation using ΔH for combustion or other cycles.
7. Gibbs Free Energy (G), Spontaneity & Equilibrium
Gibbs free energy (G): Thermodynamic potential that predicts spontaneity under constant T and P.
- ΔG = ΔH – TΔS
-
If ΔG < 0: Process is spontaneous
If ΔG = 0: Equilibrium
If ΔG > 0: Non-spontaneous
- Relates energy, entropy, and feasibility of chemical processes.
Only processes with negative free energy change (ΔG < 0) occur spontaneously at constant T, P.
- At equilibrium: ΔG° = –RT ln K (K = equilibrium constant)
8. Standard States & Thermodynamic Data
- Standard state: Most stable form of a substance at 1 bar (or 1 atm) and 298 K.
- Standard enthalpy, entropy, and Gibbs free energy values (ΔH°, S°, ΔG°) are tabulated for many substances.
- Enthalpy/entropy of elements in standard state (ΔH_f°, S°) is usually set to zero for reference.
9. Examples and Typical Numericals
-
(a) Calculating ΔH from ΔU:
2H₂(g) + O₂(g) → 2H₂O(l); ΔU = –572 kJ
ΔH = ΔU + Δn_gRT = –572 + (–2)×8.314×298/1000 ≈ –577 kJ
-
(b) Use Hess’s Law:
If C(s) + O₂(g) → CO₂(g); ΔH₁
C(s) + ½O₂(g) → CO(g); ΔH₂
CO(g) + ½O₂(g) → CO₂(g); ΔH₃
Then ΔH₁ = ΔH₂ + ΔH₃
-
(c) Spontaneity:
If ΔH = +20 kJ, ΔS = +0.150 kJ/K at 300 K:
ΔG = 20 – 300×0.150 = –25 kJ; Spontaneous
1. Introduction & Basic Concepts
Thermodynamics is the branch of science that deals with the quantitative relationship between heat and other forms of energy during physical and chemical changes.
- System: Part of universe under study (e.g., a reaction vessel).
- Surroundings: Everything outside the system.
- Boundary: Separates system from surroundings—can be real or imaginary.
-
Types of Systems:
- Open: Exchange mass & energy (e.g., open beaker).
- Closed: Exchange energy only (e.g., stoppered flask).
- Isolated: No exchange (e.g., thermos flask).
- State Functions: Depend only on initial & final state (e.g., internal energy U, enthalpy H, entropy S).
- Path Functions: Depend on path/process (e.g., heat q, work w).
Knowing which quantities are state or path functions is crucial for numericals and MCQs!
2. Intensive and Extensive Properties
Intensive properties do not depend on the amount of substance present (e.g., temperature, pressure, density, boiling point).
Extensive properties depend on the quantity of matter present (e.g., mass, volume, internal energy, enthalpy).
- Intensive Property: Value is same for any sample size.
- Examples: Temperature, pressure, refractive index, density, boiling point, concentration
- Extensive Property: Value changes with amount of substance.
- Examples: Mass, volume, energy, enthalpy, entropy, heat capacity
| Property |
Intensive |
Extensive |
| Temperature |
✓ |
|
| Volume |
|
✓ |
| Density |
✓ |
|
| Mass |
|
✓ |
| Enthalpy |
|
✓ |
| Molarity |
✓ |
|
| Entropy |
|
✓ |
| Heat capacity |
|
✓ |
Intensive properties are useful in identifying substances; extensive properties are additive for subsystems.
30 Thermodynamics MCQs – Practice Test
| MCQ | Options |
|---|
| 1. Thermodynamics deals with: | a) rate of reaction b) energy changes c) atom structure d) all above |
| 2. Which is a state function? | a) Work b) Heat c) Enthalpy d) Path |
| 3. The internal energy of an isolated system: | a) increases b) decreases c) changes randomly d) remains constant |
| 4. Work done during isothermal reversible expansion of an ideal gas is: | a) zero b) minimum c) maximum d) none |
| 5. The SI unit of entropy is: | a) J K–1mol–1 b) J mol–1 c) kJ d) none |
| 6. ΔH = ΔU + ΔnRT is applicable for: | a) solids b) liquids c) gases d) all |
| 7. For a spontaneous process, ΔG is: | a) > 0 b) < 0 c) 0 d) positive |
| 8. Which of the following is an extensive property? | a) Pressure b) Temperature c) Volume d) Density |
| 9. If heat is absorbed by system, q is: | a) +ve b) –ve c) zero d) undefined |
| 10. Standard enthalpy of formation for elements in standard state: | a) 1 b) 0 c) >0 d) <0 |
| 11. The enthalpy change for neutralization of strong acid and strong base: | a) always positive b) always zero c) always negative d) always equal to –57.1 kJ |
| 12. Which law states energy can neither be created nor destroyed? | a) First b) Second c) Third d) Zeroth |
| 13. If a reaction has ΔG > 0 it is: | a) spontaneous b) at equilibrium c) non-spontaneous d) exothermic |
| 14. For a process at equilibrium, ΔG is: | a) < 0 b) > 0 c) 0 d) undefined |
| 15. The value of ΔH for exothermic reactions is: | a) positive b) negative c) zero d) undefined |
| 16. The total entropy change of universe is: | a) always zero b) always negative c) always positive for spontaneous processes d) undefined |
| 17. Enthalpy of sublimation equals: | a) ΔH_vap + ΔH_fus b) ΔH_fus c) ΔH_vap d) None |
| 18. In which process is q = 0? | a) Isothermal b) Isochoric c) Adiabatic d) Isobaric |
| 19. For isochoric process, work done is: | a) maximum b) minimum c) zero d) negative |
| 20. For 1st law, ΔU = q + w; which is true for work done by system? | a) w is +ve b) w is –ve c) w = 0 d) w = q |
| 21. If the sign of ΔG is negative, the reaction is: | a) reversible b) exothermic c) spontaneous d) non-spontaneous |
| 22. Entropy change for freezing of water: | a) +ve b) –ve c) zero d) maximum |
| 23. Which of the following is NOT a thermodynamic state function? | a) Enthalpy b) Internal energy c) Entropy d) Heat |
| 24. First law fails to explain: | a) direction b) feasibility c) both a & b d) none |
| 25. Lattice enthalpy is highest for: | a) NaCl b) KCl c) MgO d) CaO |
| 26. The SI unit of enthalpy is: | a) cal b) kJ c) J d) kJ mol–1 |
| 27. Standard state for solids is: | a) amorphous b) crystalline c) liquid d) gas |
| 28. Which equation is the mathematical statement of 1st law? | a) ΔU = q + w b) ΔS = q/T c) ΔG = ΔH – TΔS d) ΔH = ΔU + PΔV |
| 29. ΔH, ΔS, ΔG in SI: | a) cal b) kJ c) J d) J or kJ/mol |
| 30. Spontaneous process is always: | a) fast b) slow c) rapid d) feasible thermodynamically |
MCQ Answers & Explanation
- b) energy changes Thermodynamics is the study of energy changes.
- c) Enthalpy State functions: U, H, S, G, not work/heat.
- d) remains constant Internal energy conserved in isolated system.
- c) maximum Work is maximum in reversible isothermal expansion.
- a) J K–1mol–1 Entropy SI: J K–1mol–1.
- c) gases Applies for gases: ΔH = ΔU + ΔnRT
- b) < 0 Spontaneous: ΔG < 0
- c) Volume Volume is extensive (depends on amount).
- a) +ve If heat is absorbed by system, q is positive.
- b) 0 Enthalpy of element in standard state is zero.
- d) always equal to –57.1 kJ For strong acid-strong base: –57.1 kJ/mol.
- a) First First law: energy can’t be created/destroyed.
- c) non-spontaneous ΔG > 0: not spontaneous.
- c) 0 At equilibrium, ΔG = 0.
- b) negative Exothermic: ΔH < 0.
- c) always positive for spontaneous processes Second law: entropy of universe increases.
- a) ΔH_vap + ΔH_fus Sublimation = vaporization + fusion.
- c) Adiabatic q = 0 for adiabatic processes.
- c) zero Isochoric: ΔV = 0, so w = 0.
- b) w is –ve w is negative for work done by system.
- c) spontaneous ΔG < 0 = spontaneous.
- b) –ve Freezing: order increases, S decreases.
- d) Heat Heat is a path function, not state function.
- c) both a & b First law cannot predict direction/feasibility.
- c) MgO MgO has highest lattice energy (high charge, small size).
- d) kJ mol–1 Standard enthalpy in kJ/mol.
- b) crystalline Standard state for solids is crystalline.
- a) ΔU = q + w This is the mathematical statement of 1st law.
- d) J or kJ/mol ΔH, ΔS, ΔG: SI J or kJ/mol.
- d) feasible thermodynamically Spontaneous = feasible, not necessarily fast.