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Thermodynamics – Class 11 Chemistry notes, important questions and MCQs

NCERT Class 11 Chemistry, chapter 5. Notes, key terms, important questions with answers and MCQs.

Chapter overview

Chemical thermodynamics deals with energy changes in reactions. Enthalpy, entropy and Gibbs energy tell how much heat is involved and whether a reaction is spontaneous.

Enthalpy

First law: ΔU = q + w; at constant pressure, ΔH = ΔU + ΔnRT.
Exothermic ΔH < 0; endothermic ΔH > 0.
Hess's law: total enthalpy change is the same whatever the path; standard enthalpies of formation, combustion and bond enthalpies.

Spontaneity

Entropy S measures disorder; for spontaneous change ΔS_total > 0.
Gibbs energy ΔG = ΔH − TΔS; ΔG < 0 spontaneous, ΔG = 0 equilibrium.
ΔG° = −RT ln K = −2.303RT log K.

Key terms

Enthalpy: heat content at constant pressure, H = U + PV
Entropy: measure of randomness of a system
Standard enthalpy of formation: enthalpy change when 1 mol of a compound forms from its elements in their standard states

Formulas

ΔH = ΔU + ΔnRT
ΔG = ΔH − TΔS
ΔG° = −2.303RT log K

Important questions with answers

What is the difference between extensive and intensive properties? Give two examples of each.2 marks

Extensive properties depend on the amount of matter (mass, volume, enthalpy). Intensive properties do not (temperature, pressure, density).

Calculate ΔH for the reaction CH₄ + 2O₂ → CO₂ + 2H₂O given ΔfH°: CH₄ = −74.8, CO₂ = −393.5, H₂O(l) = −285.8 kJ/mol.3 marks

ΔH = [−393.5 + 2(−285.8)] − [−74.8 + 0] = −965.1 + 74.8 = −890.3 kJ/mol.

Explain the conditions of spontaneity using ΔG = ΔH − TΔS for the four combinations of signs of ΔH and ΔS, with one example each.5 marks

(1) ΔH < 0, ΔS > 0: ΔG always negative, spontaneous at all temperatures (combustion of fuels). (2) ΔH > 0, ΔS < 0: ΔG always positive, never spontaneous (making ozone from oxygen). (3) ΔH < 0, ΔS < 0: spontaneous at low temperature, when |ΔH| > |TΔS| (freezing of water below 0 °C). (4) ΔH > 0, ΔS > 0: spontaneous at high temperature, when TΔS > ΔH (melting of ice above 0 °C, decomposition of CaCO₃ on strong heating).

Quick revision

Hess's law for indirect enthalpies.
ΔG < 0 means spontaneous.
Entropy of gases > liquids > solids.

Exam tips

ΔH < 0 and ΔS > 0 ⇒ spontaneous at all temperatures.
ΔH°reaction = Σ ΔfH°(products) − Σ ΔfH°(reactants); ΔfH° of elements in standard states is zero.

Practice MCQs

A reaction that releases heat is
1. endothermic
2. adiabatic
3. isothermal
4. exothermic
Show answer
Answer: (d) exothermic. ΔH < 0.
A reaction is spontaneous when ΔG is
1. positive
2. zero
3. negative
4. infinite
Show answer
Answer: (c) negative. ΔG < 0.
Entropy increases in
1. freezing of water
2. formation of a crystal
3. condensation of steam
4. melting of ice
Show answer
Answer: (d) melting of ice. Solid → liquid.
Hess's law is based on
1. conservation of mass
2. entropy
3. enthalpy being a state function
4. equilibrium
Show answer
Answer: (c) enthalpy being a state function. Path-independent.
For an isolated system, which is always true
1. ΔG > 0
2. ΔH = 0
3. ΔS < 0
4. ΔU = 0
Show answer
Answer: (d) ΔU = 0. No exchange of heat or work.
A reaction with ΔH = +40 kJ and ΔS = +100 J/K becomes spontaneous above
1. 4000 K
2. 40 K
3. 400 K
4. 100 K
Show answer
Answer: (c) 400 K. T > ΔH/ΔS.
For N₂(g) + 3H₂(g) → 2NH₃(g), Δn is
1. +2
2. −2
3. 0
4. −4
Show answer
Answer: (b) −2. 2 − 4.
For the reaction C(s) + O₂(g) → CO₂(g), ΔH = −393.5 kJ. Heat released on burning 6 g of carbon is
1. 393.5 kJ
2. 787 kJ
3. 196.75 kJ
4. 98.4 kJ
Show answer
Answer: (c) 196.75 kJ. 0.5 mol.

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Written by Kundali IQ from standard, well-established facts to help with revision. This is not an official NCERT publication; always read your textbook too.