Thermodynamics and Chemical Energetics -MCQs

Thermodynamics and Chemical Energetics – 40 MCQs for NEET & JEE Main

This section provides 40 multiple-choice questions covering key concepts from Thermodynamics and Chemical Energetics, essential for your NEET and JEE Main preparation. Each question has four options, with only one correct answer.

Section 1: Multiple Choice Questions (MCQs)

Instructions: Choose the single best answer for each question.

  1. Which of the following is an intensive property? A) Volume B) Mass C) Temperature D) Internal Energy
  2. A system that can exchange energy but not matter with its surroundings is called a: A) Open system B) Closed system C) Isolated system D) Adiabatic system
  3. Which of the following is NOT a state function? A) Enthalpy (H) B) Internal Energy (U) C) Work (w) D) Gibbs Free Energy (G)
  4. For an isothermal process in an ideal gas, which of the following is true? A) ΔT=0,ΔU=0 B) ΔP=0,q=0 C) ΔV=0,w=0 D) ΔS=0,ΔG=0
  5. According to the First Law of Thermodynamics, for an isolated system: A) ΔU=q B) ΔU=w C) ΔU=0 D) q=w
  6. In an exothermic reaction, ΔH is: A) Positive B) Negative C) Zero D) Dependent on the path
  7. Which of the following processes is always spontaneous at all temperatures? A) ΔH>0,ΔS<0 B) ΔH<0,ΔS>0 C) ΔH>0,ΔS>0 D) ΔH<0,ΔS<0
  8. For the reaction N2(g)+3H2(g)→2NH3(g), what is the relationship between ΔH and ΔU? A) ΔH=ΔU B) ΔH=ΔU−2RT C) ΔH=ΔU+2RT D) ΔH=ΔU+RT
  9. The standard enthalpy of formation (ΔHf∘​) of an element in its most stable state is: A) Always positive B) Always negative C) Zero D) Undefined
  10. Hess’s Law of Constant Heat Summation is based on: A) Law of conservation of mass B) Law of conservation of energy C) Law of definite proportions D) Law of multiple proportions
  11. If ΔG∘ for a reaction is positive, the reaction is: A) Spontaneous B) Non-spontaneous C) At equilibrium D) Exothermic
  12. The entropy of the universe tends towards a maximum. This statement is related to: A) First Law of Thermodynamics B) Second Law of Thermodynamics C) Third Law of Thermodynamics D) Zeroth Law of Thermodynamics
  13. For a spontaneous process, the total entropy change of the universe (ΔSuniverse​) is: A) ΔSuniverse​<0 B) ΔSuniverse​=0 C) ΔSuniverse​>0 D) ΔSuniverse​ can be positive, negative, or zero
  14. Which of the following processes leads to an increase in entropy? A) Freezing of water B) Condensation of steam C) Sublimation of dry ice D) Cooling a gas
  15. The formula for calculating work done during irreversible expansion of a gas is: A) w=Pext​ΔV B) w=−Pext​ΔV C) w=−nRTln(V2​/V1​) D) w=ΔU−q
  16. For a reversible isothermal expansion of an ideal gas, the work done is maximum. What is ΔU for this process? A) Positive B) Negative C) Zero D) Dependent on pressure
  17. The Born-Haber cycle is used to calculate: A) Bond enthalpy B) Enthalpy of neutralization C) Lattice energy D) Electron gain enthalpy
  18. Which statement is CORRECT for an endothermic reaction? A) Heat is released to the surroundings. B) ΔH is negative. C) The products have higher enthalpy than reactants. D) The reaction is always spontaneous.
  19. If ΔH=−100 kJ and ΔS=−200 J K−1 for a reaction, at what temperature will the reaction become spontaneous? A) Below 500 K B) Above 500 K C) At 500 K D) The reaction is never spontaneous
  20. The Third Law of Thermodynamics states that the entropy of a perfectly crystalline substance at 0 K is: A) Positive B) Negative C) Zero D) Undefined
  21. What are the signs of ΔH and ΔS for the melting of ice at 0°C? A) ΔH>0,ΔS>0 B) ΔH<0,ΔS<0 C) ΔH>0,ΔS<0 D) ΔH<0,ΔS>0
  22. For a reaction where ΔH=+20 kJ and ΔS=+50 J K−1, at what temperature will the reaction be at equilibrium? A) 200 K B) 300 K C) 400 K D) 500 K
  23. A process is non-spontaneous if: A) ΔG<0 B) ΔG>0 C) ΔG=0 D) ΔSuniverse​>0
  24. The relationship between ΔG∘ and the equilibrium constant K is given by: A) ΔG∘=RTlnK B) ΔG∘=−RTlnK C) ΔG∘=K−RT D) ΔG∘=−RT/K
  25. The heat absorbed by the system at constant volume is equal to: A) ΔH B) ΔU C) w D) ΔG
  26. Which of the following is an example of an isolated system? A) A cup of hot coffee on a table B) A sealed reaction vessel in a laboratory C) Water in a well-insulated thermos flask D) A balloon filled with air
  27. If work is done BY the system, according to IUPAC convention, ‘w’ is: A) Positive B) Negative C) Zero D) Undefined
  28. For the reaction C(s)+O2(g)→CO2(g), which relation is approximately correct? A) ΔH=ΔU B) ΔH>ΔU C) ΔH<ΔU D) ΔH=ΔU+RT
  29. The standard enthalpy of formation of C(graphite) is: A) Positive B) Negative C) Zero D) Cannot be determined
  30. Which of the following will have the highest value of standard enthalpy of atomization? A) Cl2(g) B) O2(g) C) N2(g) D) H2(g)
  31. If ΔG∘=0 for a reaction, then: A) K = 0 B) K = 1 C) K > 1 D) K < 1
  32. The process of diffusion of gases in an isolated container is: A) Spontaneous and reversible B) Spontaneous and irreversible C) Non-spontaneous and reversible D) Non-spontaneous and irreversible
  33. Which of the following factors does NOT affect the entropy of a system? A) Temperature B) Volume (for gases) C) Number of moles of gas D) Path taken for the process
  34. For a reaction where ΔH<0 and ΔS<0, the reaction is spontaneous: A) At all temperatures B) At high temperatures C) At low temperatures D) Never spontaneous
  35. The concept of Born-Haber cycle is typically applied to calculate lattice energy for: A) Covalent compounds B) Molecular compounds C) Ionic compounds D) Metallic compounds
  36. Which quantity is a measure of the disorder of a system? A) Enthalpy B) Internal energy C) Gibbs free energy D) Entropy
  37. If q is negative, it means: A) Heat is absorbed by the system. B) Heat is released by the system. C) Work is done on the system. D) Work is done by the system.
  38. For a cyclic process, the net change in internal energy (ΔU) is: A) Positive B) Negative C) Zero D) Dependent on the path
  39. The relation ΔH=ΔU+Δng​RT is applicable to: A) Solids only B) Liquids only C) Gases only D) Reactions involving gases
  40. Which of the following is true for a reversible process? A) Occurs rapidly. B) ΔSuniverse​>0. C) System and surroundings are always in equilibrium. D) Work obtained is minimum.

Section 2: Answer Key

  1. C
  2. B
  3. C
  4. A
  5. C
  6. B
  7. B
  8. B
  9. C
  10. B
  11. B
  12. B
  13. C
  14. C
  15. B
  16. C
  17. C
  18. C
  19. A
  20. C
  21. A
  22. C
  23. B
  24. B
  25. B
  26. C
  27. B
  28. A
  29. C
  30. C
  31. B
  32. B
  33. D
  34. C
  35. C
  36. D
  37. B
  38. C
  39. D
  40. C

Section 3: Detailed Explanations

  1. C) Temperature
    • Explanation: Intensive properties do not depend on the amount of matter. Temperature, pressure, density are intensive. Volume, mass, and internal energy are extensive as they depend on the amount of substance.
  2. B) Closed system
    • Explanation: A closed system can exchange energy (heat and work) but not matter with its surroundings. An open system exchanges both, and an isolated system exchanges neither.
  3. C) Work (w)
    • Explanation: Work (w) and Heat (q) are path functions, meaning their values depend on the specific path taken between initial and final states. Internal Energy (U), Enthalpy (H), Entropy (S), and Gibbs Free Energy (G) are state functions, depending only on the initial and final states.
  4. A) ΔT=0,ΔU=0
    • Explanation: For an isothermal process, the temperature (T) remains constant. For an ideal gas, internal energy (U) depends only on temperature, so if T is constant, ΔU=0.
  5. C) ΔU=0
    • Explanation: The First Law of Thermodynamics states ΔU=q+w. For an isolated system, there is no exchange of heat (q=0) and no exchange of work (w=0) with the surroundings. Therefore, ΔU for an isolated system is zero.
  6. B) Negative
    • Explanation: Exothermic reactions release heat to the surroundings, meaning the enthalpy of the system decreases. By convention, a decrease in enthalpy is represented by a negative ΔH value.
  7. B) ΔH<0,ΔS>0
    • Explanation: For a process to be spontaneous, ΔG=ΔH−TΔS must be negative. If ΔH is negative (exothermic) and ΔS is positive (increase in disorder), then the −TΔS term will always be negative, making ΔG always negative, irrespective of temperature.
  8. B) ΔH=ΔU−2RT
    • Explanation: The relationship is ΔH=ΔU+Δng​RT. For N2(g)+3H2(g)→2NH3(g): Moles of gaseous products (Δng​ products) = 2 Moles of gaseous reactants (Δng​ reactants) = 1 + 3 = 4 Δng​ = 2 – 4 = -2. Therefore, ΔH=ΔU+(−2)RT=ΔU−2RT.
  9. C) Zero
    • Explanation: By convention, the standard enthalpy of formation (ΔHf∘​) of an element in its most stable elemental form (e.g., O2(g), C(graphite), Na(s)) under standard conditions is taken as zero.
  10. B) Law of conservation of energy
    • Explanation: Hess’s Law states that the total enthalpy change for a reaction is independent of the pathway taken (number of steps). This is a direct consequence of the First Law of Thermodynamics (conservation of energy), as enthalpy is a state function.
  11. B) Non-spontaneous
    • Explanation: The criterion for spontaneity at constant temperature and pressure is based on Gibbs Free Energy. If ΔG∘>0, the reaction is non-spontaneous in the forward direction under standard conditions.
  12. B) Second Law of Thermodynamics
    • Explanation: The Second Law of Thermodynamics states that for any spontaneous process, the entropy of the universe always increases (ΔSuniverse​>0). This implies a natural tendency towards greater disorder.
  13. C) ΔSuniverse​>0
    • Explanation: The Second Law of Thermodynamics explicitly states that for a spontaneous process, the total entropy of the universe (system + surroundings) must increase.
  14. C) Sublimation of dry ice
    • Explanation: Entropy is a measure of disorder.
      • Freezing (liquid to solid) and condensation (gas to liquid) decrease disorder, so entropy decreases.
      • Cooling a gas decreases the kinetic energy and thus disorder.
      • Sublimation (solid to gas) involves a transition to a much more disordered gaseous state, so entropy increases significantly.
  15. B) w=−Pext​ΔV
    • Explanation: For an irreversible process (e.g., expansion against a constant external pressure), the work done is given by w=−Pext​ΔV. The negative sign indicates work done by the system during expansion.
  16. C) Zero
    • Explanation: For an ideal gas, internal energy (U) is solely a function of temperature. In an isothermal process, the temperature is constant (ΔT=0), so there is no change in internal energy, i.e., ΔU=0.
  17. C) Lattice energy
    • Explanation: The Born-Haber cycle is a thermochemical cycle that applies Hess’s Law to calculate the lattice energy of an ionic compound indirectly by relating it to other measurable enthalpy changes (e.g., ionization energy, electron gain enthalpy, enthalpy of formation, atomization enthalpies).
  18. C) The products have higher enthalpy than reactants.
    • Explanation: For an endothermic reaction, heat is absorbed by the system, meaning the enthalpy of the products is greater than the enthalpy of the reactants. This results in a positive ΔH.
  19. A) Below 500 K
    • Explanation: For spontaneity, ΔG<0. We use ΔG=ΔH−TΔS. Given ΔH=−100 kJ=−100000 J. ΔS=−200 J K−1. For ΔG=0, T=ΔH/ΔS=(−100000 J)/(−200 J K−1)=500 K. Since ΔH is negative and ΔS is negative, the reaction is spontaneous at low temperatures (when the −TΔS term is smaller in magnitude than ΔH). Thus, spontaneous below 500 K.
  20. C) Zero
    • Explanation: The Third Law of Thermodynamics states that the entropy of a perfectly crystalline substance at absolute zero (0 K) is exactly zero. This represents a state of perfect order.
  21. A) ΔH>0,ΔS>0
    • Explanation: Melting of ice is an endothermic process (requires heat), so ΔH>0. It also involves a transition from a more ordered solid state to a less ordered liquid state, so entropy increases, ΔS>0.
  22. C) 400 K
    • Explanation: At equilibrium, ΔG=0. So, ΔH−TΔS=0⇒T=ΔH/ΔS. Given ΔH=+20 kJ=+20000 J. ΔS=+50 J K−1. T=(20000 J)/(50 J K−1)=400 K.
  23. B) ΔG>0
    • Explanation: A process is non-spontaneous in the forward direction if the change in Gibbs Free Energy (ΔG) is positive.
  24. B) ΔG∘=−RTlnK
    • Explanation: This fundamental equation relates the standard Gibbs free energy change (ΔG∘) to the equilibrium constant (K) of a reaction.
  25. B) ΔU
    • Explanation: According to the First Law, ΔU=q+w. For an isochoric process, ΔV=0, so work done (w = −Pext​ΔV) is zero. Therefore, ΔU=qV​, meaning the heat exchanged at constant volume is equal to the change in internal energy.
  26. C) Water in a well-insulated thermos flask
    • Explanation: An isolated system cannot exchange either energy (heat and work) or matter with its surroundings. A well-insulated thermos flask approximates an isolated system for a short period.
  27. B) Negative
    • Explanation: According to IUPAC sign conventions, if work is done BY the system (e.g., expansion), the system loses energy, so ‘w’ is negative. If work is done ON the system (e.g., compression), the system gains energy, so ‘w’ is positive.
  28. A) ΔH=ΔU
    • Explanation: For the reaction C(s)+O2(g)→CO2(g): Moles of gaseous products (Δng​ products) = 1 Moles of gaseous reactants (Δng​ reactants) = 1 (C is solid, so not included in Δng​) Δng​ = 1 – 1 = 0. Since Δng​=0, then ΔH=ΔU+0×RT=ΔU.
  29. C) Zero
    • Explanation: The standard enthalpy of formation of an element in its most stable reference state is defined as zero. Graphite is the most stable allotrope of carbon at standard conditions.
  30. C) N2(g)
    • Explanation: Standard enthalpy of atomization is the energy required to break all bonds in one mole of a substance to form gaseous atoms.
      • H2(g): H-H single bond
      • Cl2(g): Cl-Cl single bond
      • O2(g): O=O double bond
      • N2(g): N#N triple bond A triple bond is the strongest bond and requires the most energy to break, thus N2 will have the highest enthalpy of atomization.
  31. B) K = 1
    • Explanation: The relationship is ΔG∘=−RTlnK. If ΔG∘=0, then −RTlnK=0, which implies lnK=0. This occurs when K = 1.
  32. B) Spontaneous and irreversible
    • Explanation: Diffusion of gases from a region of higher concentration to lower concentration occurs naturally without external intervention, so it’s spontaneous. It cannot be reversed by an infinitesimal change and involves a finite change, making it an irreversible process. All natural spontaneous processes are irreversible.
  33. D) Path taken for the process
    • Explanation: Entropy (S) is a state function, meaning its change depends only on the initial and final states of the system, not on the path taken. Temperature, volume (for gases), and the number of moles of gaseous species (composition) all affect the entropy of a system.
  34. C) At low temperatures
    • Explanation: For ΔG=ΔH−TΔS: If ΔH<0 (exothermic) and ΔS<0 (decrease in disorder), ΔG will be negative (spontaneous) only when the −TΔS term is less positive than the magnitude of ΔH is negative. This happens at low temperatures (when T is small). At high T, −TΔS becomes more positive, making ΔG positive.
  35. C) Ionic compounds
    • Explanation: The Born-Haber cycle is specifically designed and primarily used to calculate the lattice energy of ionic compounds. It connects the formation of an ionic solid from its elements to a series of steps involving ionization energy, electron gain enthalpy, atomization enthalpy, and lattice energy.
  36. D) Entropy
    • Explanation: Entropy (S) is the thermodynamic quantity that measures the degree of randomness or disorder in a system.
  37. B) Heat is released by the system.
    • Explanation: According to IUPAC sign convention, if q is negative, it means heat is lost by the system to the surroundings, i.e., the process is exothermic from the system’s perspective.
  38. C) Zero
    • Explanation: A cyclic process is one in which the system returns to its initial state. Since internal energy (U) is a state function, its change depends only on the initial and final states. If initial and final states are the same, ΔU=0.
  39. D) Reactions involving gases
    • Explanation: The relation ΔH=ΔU+Δng​RT specifically accounts for the work done due to the change in the number of moles of gaseous species (Δng​). Thus, it is applicable to reactions where there is a change in the amount of gas, irrespective of whether other phases are also present.
  40. C) System and surroundings are always in equilibrium.
    • Explanation: A reversible process is an idealized process that occurs infinitesimally slowly, allowing the system and surroundings to always remain in equilibrium with each other throughout the process. This leads to maximum work output. Real processes are irreversible and occur rapidly.

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