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Thermodynamics Revision

Thermodynamics (NCERT Class 11 Chemistry) - Chemca

Thermodynamics

110 Short Q&A • NCERT Class 11 Chemistry

Class 11 JEE / NEET

110 Short Q&A for Thermodynamics

    Part 1: Basic Concepts and First Law of Thermodynamics
  1. Q1: What is the System in thermodynamics?
    A: The part of the universe chosen for thermodynamic consideration.
  2. Q2: What is the Surrounding?
    A: Everything in the universe other than the system.
  3. Q3: What is an Isolated System?
    A: A system that can exchange neither energy nor matter with the surroundings.
  4. Q4: What is a Closed System?
    A: A system that can exchange energy but not matter with the surroundings.
  5. Q5: What is an Open System?
    A: A system that can exchange both energy and matter with the surroundings.
  6. Q6: Define an Isothermal Process.
    A: A process carried out at constant temperature (ΔT=0).
  7. Q7: Define an Adiabatic Process.
    A: A process where no heat is exchanged between the system and surroundings (q=0).
  8. Q8: Define an Isobaric Process.
    A: A process carried out at constant pressure (ΔP=0).
  9. Q9: Define an Isochoric Process.
    A: A process carried out at constant volume (ΔV=0).
  10. Q10: What is a Reversible Process?
    A: A process that can be reversed by an infinitesimal change and takes place in infinite steps.
  11. Q11: What are State Functions?
    A: Properties of the system that depend only on the initial and final states, not on the path taken.
  12. Q12: Give three examples of State Functions.
    A: Internal Energy (U), Enthalpy (H), Entropy (S), Gibbs Energy (G).
  13. Q13: Give two examples of Path Functions.
    A: Heat (q) and Work (w).
  14. Q14: State the First Law of Thermodynamics.
    A: Energy can neither be created nor destroyed; the total energy of the universe is constant.
  15. Q15: Write the mathematical form of the First Law of Thermodynamics.
    A: ΔU = q + w (Change in internal energy equals heat plus work).
  16. Q16: What is the sign convention for heat absorbed by the system?
    A: Positive (+q).
  17. Q17: What is the sign convention for work done on the system (compression)?
    A: Positive (+w).
  18. Q18: What is the sign convention for work done by the system (expansion)?
    A: Negative (−w).
  19. Q19: What is the term for the heat content of a system at constant pressure?
    A: Enthalpy (H).
  20. Q20: Write the relationship between ΔH and ΔU for chemical reactions.
    A: ΔH = ΔU + ΔngRT (for reactions involving gases).
    21. Part 2: Enthalpy Changes (ΔH) and Hess's Law
  21. Q21: Define Enthalpy of Reaction (ΔrH).
    A: The enthalpy change accompanying a reaction represented by a balanced equation.
  22. Q22: What is an Exothermic Reaction?
    A: A reaction where heat is released (ΔH is negative).
  23. Q23: What is an Endothermic Reaction?
    A: A reaction where heat is absorbed (ΔH is positive).
  24. Q24: Define Standard Enthalpy of Formation (ΔfH°).
    A: Enthalpy change when one mole of a compound is formed from its elements in their standard states.
  25. Q25: What is the standard enthalpy of formation of an element in its reference state?
    A: Zero (By definition, e.g., ΔfH° for C(graphite) = 0).
  26. Q26: Define Enthalpy of Combustion (ΔcH°).
    A: Enthalpy change when one mole of a substance is completely burnt in excess oxygen.
  27. Q27: Define Enthalpy of Neutralization.
    A: Enthalpy change when one gram equivalent of an acid is neutralized by one gram equivalent of a base in dilute solution.
  28. Q28: State Hess's Law of Constant Heat Summation.
    A: The total enthalpy change for a reaction is the same regardless of the path taken, provided the initial and final states are the same.
  29. Q29: What is Bond Enthalpy?
    A: The energy required to break one mole of a particular type of bond in the gaseous state.
  30. Q30: How is ΔrH calculated using bond enthalpies?
    A: ΔrH = ∑Hbonds broken − ∑Hbonds formed (Reactants - Products).
  31. Q31: How is ΔrH° calculated using standard enthalpies of formation?
    A: ΔrH° = ∑ΔfH°products − ∑ΔfH°reactants.
  32. Q32: Define Enthalpy of Phase Transition.
    A: The enthalpy change when a substance changes from one physical state to another at constant T and P.
  33. Q33: What is the relationship between ΔHsublimation and ΔHfusion/ΔHvaporization?
    A: ΔHsub = ΔHfus + ΔHvap.
  34. Q34: Define Standard State of a substance.
    A: Its most stable state at 1 bar pressure and a specified temperature (usually 298 K).
  35. Q35: What does the term Δng in ΔH = ΔU + ΔngRT represent?
    A: (Moles of gaseous products) - (Moles of gaseous reactants).
    36. Part 3: Second and Third Laws, Entropy, and Spontaneity
  36. Q36: State the Second Law of Thermodynamics (Entropy).
    A: The entropy of the universe always increases in the course of every spontaneous process.
  37. Q37: Define Entropy (S).
    A: A measure of the degree of randomness or disorder in the system.
  38. Q38: What is the unit of Entropy?
    A: J K−1 mol−1 (or J K−1).
  39. Q39: How does entropy change during the melting of ice?
    A: Increases (ΔS > 0), as liquid is more disordered than solid.
  40. Q40: How is the change in entropy for a reversible process calculated?
    A: ΔS = qrev / T.
  41. Q41: When is a process Spontaneous (based on ΔStotal)?
    A: When ΔStotal (system + surr) > 0.
  42. Q42: What is the spontaneity condition based on Gibbs Free Energy (ΔG)?
    A: Process is spontaneous if ΔG < 0.
  43. Q43: Write the equation that relates ΔG, ΔH, and ΔS (Gibbs equation).
    A: ΔG = ΔH − TΔS.
  44. Q44: What does Gibbs Free Energy (G) represent?
    A: The energy available to do useful work.
  45. Q45: What does ΔG = 0 signify for a reaction?
    A: The system is at Equilibrium.
  46. Q46: State the Third Law of Thermodynamics.
    A: The entropy of a perfectly crystalline substance at absolute zero (0 K) is zero.
  47. Q47: How do the signs of ΔH and ΔS lead to spontaneity at all temperatures?
    A: ΔH < 0 (Exothermic) and ΔS > 0 (Increased disorder).
  48. Q48: How do the signs of ΔH and ΔS lead to non-spontaneity at all temperatures?
    A: ΔH > 0 (Endothermic) and ΔS < 0 (Decreased disorder).
  49. Q49: What condition is required for a ΔH > 0, ΔS > 0 process to be spontaneous?
    A: High temperature, such that TΔS > ΔH.
  50. Q50: What condition is required for a ΔH < 0, ΔS < 0 process to be spontaneous?
    A: Low temperature, such that TΔS < ΔH.
  51. Q51: How does ΔG° relate to the Equilibrium Constant (K)?
    A: ΔG° = −RT ln K (or −2.303RT log K).
  52. Q52: If ΔG° is positive, is K > 1 or K < 1?
    A: K < 1 (Reaction favors reactants at standard state).
  53. Q53: If ΔG° is negative, is K > 1 or K < 1?
    A: K > 1 (Reaction favors products at standard state).
    54. Part 4: Work and Specific Concepts
  54. Q54: What type of work is typically considered in chemical thermodynamics?
    A: Pressure-Volume (PV) Work.
  55. Q55: What is the formula for PV work done in a reversible expansion?
    A: wrev = −∫ Pext dV (where Pext ≈ Pint).
  56. Q56: What is the formula for PV work done in an irreversible expansion (against constant external pressure)?
    A: wirrev = −PextΔV.
  57. Q57: Which type of process (reversible or irreversible) does maximum work?
    A: Reversible expansion work is maximum.
  58. Q58: What is the value of work done in a vacuum (free expansion)?
    A: Zero (w = −PextΔV, and Pext = 0).
  59. Q59: What is the relationship between the specific heat capacity at constant volume (Cv) and internal energy?
    A: Cv = (∂U / &partial;T)v.
  60. Q60: What is the relationship between the specific heat capacity at constant pressure (Cp) and enthalpy?
    A: Cp = (∂H / &partial;T)p.
  61. Q61: What is Joule-Thomson effect?
    A: The cooling or heating of a real gas when it is made to expand adiabatically from a region of high pressure to a region of low pressure.
  62. Q62: What is the relationship between Cp and Cv for an ideal gas (Mayer's formula)?
    A: Cp − Cv = R (where R is the Universal Gas Constant).
  63. Q63: Define Standard Enthalpy of Atomization (ΔaH°).
    A: Enthalpy change when one mole of a substance is completely dissociated into atoms in the gaseous state.
  64. Q64: Define Standard Enthalpy of Solution (ΔsolH°).
    A: Enthalpy change when one mole of a substance dissolves in a specified amount of solvent.
  65. Q65: What is the standard pressure taken for gas measurements in NCERT?
    A: 1 bar.
    66. Part 5: Standard Gibbs Energy and Equilibrium
  66. Q66: Define Standard Gibbs Energy of Formation (ΔfG°).
    A: Change in Gibbs energy when one mole of a compound is formed from its elements in their standard states.
  67. Q67: What is the value of ΔfG° for O2 gas?
    A: Zero (element in standard state).
  68. Q68: What is the relationship between ΔG and ΔG°?
    A: ΔG = ΔG° + RT ln Q (where Q is the reaction quotient).
  69. Q69: At equilibrium, what is the value of the reaction quotient Q?
    A: Q = K (Equilibrium Constant).
  70. Q70: If ΔH and ΔS are both positive, which factor determines spontaneity?
    A: Temperature (T) (Spontaneous at high T).
  71. Q71: If K = 1, what is the value of ΔG°?
    A: ΔG° = 0.
  72. Q72: If ΔH = −10 kJ and ΔS = −100 J/K, calculate Teq.
    A: T = ΔH / ΔS = 100 K.
  73. Q73: What is the term for a reaction that goes almost to completion?
    A: High K or Large negative ΔG°.
  74. Q74: What is the term for a reaction that hardly proceeds?
    A: Low K or Large positive ΔG°.
  75. Q75: How does increasing the temperature affect the ΔG of an exothermic reaction?
    A: ΔG = ΔH − TΔS; for ΔH < 0, increasing T makes TΔS (positive) more significant, making ΔG less negative (less spontaneous).
  76. Q76: How does increasing the temperature affect the ΔG of an endothermic reaction?
    A: ΔG = ΔH − TΔS; for ΔH > 0, increasing T makes TΔS more significant, potentially making ΔG negative (more spontaneous).
    77. Part 6: Miscellaneous Concepts and ΔU vs ΔH
  77. Q77: What is the value of ΔV for a reaction in a bomb calorimeter?
    A: Zero (constant volume).
  78. Q78: What quantity is measured directly by a bomb calorimeter?
    A: ΔU (Change in Internal Energy).
  79. Q79: What is the relationship between ΔH and ΔU when there are no gaseous reactants or products?
    A: ΔH = ΔU (Δng = 0).
  80. Q80: Is ΔH or ΔU usually larger for a gaseous reaction where the number of moles of product gas is greater than reactant gas?
    A: ΔH > ΔU (since Δng > 0, and ΔH = ΔU + positive term).
  81. Q81: What is the maximum work a system can do under isothermal and reversible conditions?
    A: wmax = −2.303nRT log(V2/V1).
  82. Q82: What is the internal energy (U) of an ideal gas dependent upon?
    A: Temperature only.
  83. Q83: What is a Cyclic Process?
    A: A process where the system's initial and final states are the same (ΔU = 0, ΔH = 0, …).
  84. Q84: What is the ΔU of a cyclic process?
    A: Zero (ΔU = 0).
  85. Q85: Is the cooling of a gas during adiabatic expansion endothermic or exothermic from the system's perspective?
    A: Neither. It is an adiabatic process (q = 0). The work done by the system decreases U, causing cooling.
  86. Q86: Define Heat Capacity (C).
    A: The heat required to raise the temperature of a given amount of substance by 1°C or 1 K.
  87. Q87: Define Molar Heat Capacity (Cm).
    A: Heat capacity for one mole of a substance.
  88. Q88: What is the main difference between Heat (q) and Internal Energy (U)?
    A: q is a path function (energy in transit); U is a state function (stored energy).
  89. Q89: What is the sign of ΔH for the process A(s) → A(g) (Sublimation)?
    A: Positive (ΔH > 0, always endothermic).
  90. Q90: What is the sign of ΔS for the reaction 2SO2(g) + O2(g) → 2SO3(g)?
    A: Negative (ΔS < 0, disorder decreases as Δng = −1).
  91. Q91: What is the standard enthalpy of combustion of a fuel a measure of?
    A: Its calorific value (how much heat it releases).
  92. Q92: What is the standard temperature used in most thermodynamic calculations?
    A: 298 K (or 25°C).
  93. Q93: Name the device used to measure the heat changes at constant volume.
    A: Bomb Calorimeter.
  94. Q94: Name the device used to measure the heat changes at constant pressure.
    A: Coffee Cup Calorimeter (or any simple calorimeter).
  95. Q95: What is the relationship between the spontaneity of a forward and reverse reaction?
    A: If the forward reaction is spontaneous, the reverse reaction is non-spontaneous (ΔGrev = −ΔGfwd).
  96. Q96: What is the ΔH value for an ideal reversible process at constant T and P?
    A: ΔH = 0 (Isothermal process for ideal gas).
  97. Q97: What is the significance of the fact that diamond is not the standard state of carbon?
    A: Its ΔfH° is not zero (graphite is the standard state).
  98. Q98: Is ΔSsystem positive or negative for the process CaCO3(s) → CaO(s) + CO2(g)?
    A: Positive (ΔS > 0), due to the formation of a gas.
  99. Q99: How does the kinetic energy of ideal gas molecules change during isothermal expansion?
    A: Remains unchanged (KE depends only on T).
  100. Q100: What is the value of ΔG when a reaction mixture is composed only of elements in their standard states?
    A: ΔG is the ΔrG for the formation of the product.
  101. Q101: What term refers to a state that is stable under a given set of conditions, but not the most stable?
    A: Metastable state (e.g., diamond at 298 K, 1 bar).
  102. Q102: What is the sign of w for a reaction in a closed container where P increases?
    A: Zero (w=0 as ΔV=0).
  103. Q103: What is the TΔS term in the Gibbs equation known as?
    A: Unavailable energy (energy lost to disorder/randomness).
  104. Q104: What is the ΔU for the isothermal expansion of an ideal gas?
    A: Zero (ΔU=0).
  105. Q105: If a system does work on the surrounding and releases heat, what is the sign of ΔU?
    A: Negative (ΔU = q + w, where q < 0 and w < 0).
  106. Q106: Why are the Cp and Cv values for liquids and solids nearly the same?
    A: They are almost incompressible, so ΔV ≈ 0 even at constant P, making PΔV ≈ 0.
  107. Q107: What is the criterion for a process to be non-spontaneous?
    A: ΔG > 0 or ΔStotal < 0.
  108. Q108: What is the ΔH for the formation of one mole of Cl(g) from Cl2(g)?
    A: 1/2 × Bond Enthalpy of Cl−Cl.
  109. Q109: Is work always zero in an isochoric process?
    A: Yes, because ΔV = 0, so w = −PextΔV = 0.
  110. Q110: If a process is spontaneous, is it necessarily fast?
    A: No, spontaneity is a thermodynamic concept; speed is a kinetic concept.

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