Take 5 chapter-wise practice tests of 25 questions each on Thermodynamics for NEET with +4/-1 scoring, answer review, and concise explanations.
5 original practice tests, 25 questions each, NEET 4/-1 marking, and answer review after submission.
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1. A system that can exchange both matter and energy with surroundings is:
Explanation: Open systems exchange both matter and energy.
2. A system that exchanges energy but not matter is:
Explanation: Closed systems do not exchange mass.
3. A system that exchanges neither matter nor energy is:
Explanation: It is ideally cut off from surroundings.
4. Internal energy is a:
Explanation: Its change depends only on initial and final states.
5. Work is a:
Explanation: Its value depends on the process path.
6. Heat is a:
Explanation: Heat depends on the process, not just state.
7. The first law of thermodynamics is:
Explanation: With chemistry sign convention, work done on system is positive.
8. In chemistry sign convention, expansion work done by the system is:
Explanation: Energy leaves the system as it does work on surroundings.
9. Compression work done on the system is:
Explanation: Work enters the system from surroundings.
10. At constant volume, heat exchanged equals change in:
Explanation: Since no PV work is done, $q_v = \Delta U$.
11. At constant pressure, heat exchanged equals change in:
Explanation: For only PV work, $q_p = \Delta H$.
12. Enthalpy is defined as:
Explanation: Enthalpy combines internal energy with pressure-volume term.
13. In an isothermal process for an ideal gas, temperature remains:
Explanation: Isothermal literally means constant temperature.
14. In an adiabatic process, heat exchange q is:
Explanation: No heat enters or leaves the system.
15. For irreversible expansion against constant external pressure, work is:
Explanation: This is the chemistry sign convention for expansion work.
16. If a gas expands from 2 L to 5 L against 1 atm external pressure, the sign of work is:
Explanation: Expansion means the system does work on surroundings.
17. A cyclic process has net change in internal energy equal to:
Explanation: State function returns to its initial value over a cycle.
18. 1 L atm is approximately equal to:
Explanation: This conversion is useful in work calculations.
19. Heat capacity is the heat required to raise the temperature of a system by:
Explanation: That is the basic definition.
20. Molar heat capacity refers to heat needed per:
Explanation: It is a molar-scale heat capacity.
21. An exothermic process has enthalpy change:
Explanation: Heat is released to surroundings.
22. An endothermic process has enthalpy change:
Explanation: Heat is absorbed by the system.
23. The value of $\Delta U$ for a process depends on:
Explanation: Internal energy is a state function.
24. The first law is fundamentally a statement of conservation of:
Explanation: Energy is neither created nor destroyed.
25. Thermodynamics begins with the most useful distinction between:
Explanation: This distinction clears up most conceptual confusion early.
26. Standard enthalpy of formation refers to formation of 1 mol compound from its elements in:
Explanation: That is the definition of standard enthalpy of formation.
27. Standard enthalpy of formation of an element in its standard state is:
Explanation: This convention simplifies enthalpy calculations.
28. Hess's law is based on the fact that enthalpy is a:
Explanation: So total enthalpy change is path independent.
29. If a reaction occurs in several steps, total enthalpy change is:
Explanation: Hess's law permits algebraic addition.
30. Bond dissociation enthalpy is usually:
Explanation: Energy is required to break a bond.
31. Enthalpy of atomisation is the enthalpy change for:
Explanation: It converts the element to isolated gaseous atoms.
32. Combustion reactions are generally:
Explanation: Burning releases substantial heat.
33. Enthalpy of neutralisation for strong acid-strong base is close to:
Explanation: It corresponds to formation of one mole of water from H$^+$ and OH$^-$.
34. Dissolution of ammonium nitrate in water is:
Explanation: Its temperature-lowering effect reveals positive enthalpy change.
35. Lattice enthalpy is associated with:
Explanation: It measures ionic crystal stability energetically.
36. Born-Haber cycle is mainly used for:
Explanation: It applies Hess's law to ionic solid formation.
37. The calorific value of a fuel is related to its enthalpy of:
Explanation: More exothermic combustion gives higher energy output.
38. If bonds broken require 800 kJ and bonds formed release 1000 kJ, $\Delta H$ is:
Explanation: $\Delta H = \Sigma$ bond broken - $\Sigma$ bond formed.
39. Why are average bond enthalpies used?
Explanation: Different molecular surroundings slightly change bond strength.
40. Heat evolved in coffee-cup calorimeter at constant pressure gives:
Explanation: Open calorimeters usually operate at constant pressure.
41. Heat absorbed by calorimeter is numerically equal to heat:
Explanation: Energy conservation connects the two.
42. If a reaction equation is reversed, the sign of enthalpy change:
Explanation: Forward and reverse processes have opposite enthalpy changes.
43. If a reaction is multiplied by 2, its $\Delta H$ becomes:
Explanation: Enthalpy is an extensive quantity.
44. Enthalpy of vaporisation is always:
Explanation: Heat is needed to separate liquid molecules into vapour.
45. Enthalpy of condensation is:
Explanation: Condensation is the reverse process.
46. A reaction may be spontaneous yet endothermic because spontaneity depends on more than:
Explanation: Spontaneity involves free-energy considerations, not just heat change.
47. The strongest reason Hess's law works is that enthalpy is:
Explanation: State functions permit algebraic route changes.
48. The bond-enthalpy approximation for reaction enthalpy is:
Explanation: Breaking requires energy; forming releases energy.
49. A more negative $\Delta H$ generally means the reaction is:
Explanation: Larger heat release makes enthalpy more negative.
50. Most numerical thermochemistry questions can be solved by combining sign convention with:
Explanation: That combination handles most NEET problems in this part.
51. Entropy is commonly associated with:
Explanation: It measures dispersal of energy and matter.
52. For a spontaneous process, total entropy of universe:
Explanation: That is the entropy criterion for spontaneity.
53. Melting of ice at 0$^\circ$C increases entropy because:
Explanation: Particles gain freedom of movement.
54. Condensation of steam decreases entropy of the system because:
Explanation: Gases have the highest disorder among common states.
55. The Gibbs free energy relation is:
Explanation: This connects enthalpy and entropy effects.
56. A process is spontaneous at constant T and P if:
Explanation: Negative free energy favors spontaneity.
57. At equilibrium, $\Delta G$ is:
Explanation: No net driving force remains at equilibrium.
58. If $\Delta H 0$, the process is:
Explanation: Both terms favor a negative $\Delta G$.
59. If $\Delta H > 0$ and $\Delta S
Explanation: Both terms make $\Delta G$ positive.
60. If $\Delta H
Explanation: The negative enthalpy term dominates when T is small.
61. If $\Delta H > 0$ and $\Delta S > 0$, spontaneity is favored at:
Explanation: The positive entropy term can dominate at large T.
62. Which process increases entropy most?
Explanation: Gas phase has much greater randomness.
63. Dissolution of a crystalline ionic solid generally tends to:
Explanation: Ordered crystal breaks into dispersed ions in solution.
64. The second law of thermodynamics can be stated as:
Explanation: This gives directionality to natural processes.
65. Maximum useful work obtainable from a process at constant T and P is related to decrease in:
Explanation: $-\Delta G$ represents the maximum non-expansion work.
66. No heat engine can be 100% efficient because:
Explanation: Second-law restrictions limit conversion of heat into work.
67. When gas expands freely into vacuum, entropy of the gas:
Explanation: The gas occupies more accessible microstates.
68. Entropy of a perfect crystal at absolute zero is taken as:
Explanation: This is the third law reference point.
69. The factor $T\Delta S$ in Gibbs equation becomes more important at:
Explanation: Temperature magnifies the entropy contribution.
70. A negative $\Delta G$ indicates the process can proceed:
Explanation: Free energy acts as the driving criterion.
71. If a process has positive $\Delta G$, then the reverse process has:
Explanation: Reversing the process changes the sign of free-energy change.
72. Which has highest entropy?
Explanation: Gas phase has maximum disorder among these.
73. Which has lowest entropy?
Explanation: Third law assigns it zero entropy.
74. Spontaneous does not necessarily mean:
Explanation: Rate and spontaneity are different ideas.
75. The cleanest way to think about spontaneity is that nature balances the tendency to lower enthalpy with the tendency to:
Explanation: Gibbs free energy combines these two tendencies.
76. If a system absorbs 100 J heat and does 40 J work, change in internal energy is:
Explanation: $\Delta U=q+w$, and here w = -40 J for work done by system, so 100 - 40 = 60 J.
77. If 50 J work is done on a system and 20 J heat is lost, $\Delta U$ is:
Explanation: Here q = -20 J and w = +50 J, so $\Delta U = 30$ J.
78. If 200 J raises temperature of a sample by 10 K, heat capacity is:
Explanation: Heat capacity = q/$\Delta T$.
79. If 100 g substance with specific heat 0.5 J g$^{-1}$ K$^{-1}$ is heated by 20 K, heat absorbed is:
Explanation: $q=ms\Delta T=100\times0.5\times20$.
80. At constant pressure, if a reaction releases 40 kJ heat, $\Delta H$ is:
Explanation: Heat release at constant pressure means negative enthalpy change.
81. At constant volume, release of 25 kJ heat implies $\Delta U$ is:
Explanation: At constant volume, $q_v = \Delta U$.
82. If reaction A has $\Delta H = -100$ kJ and reaction B has $\Delta H = +40$ kJ, combined enthalpy is:
Explanation: Add algebraically.
83. If total bond enthalpy of reactants is 1500 kJ and of products is 1700 kJ, reaction enthalpy is:
Explanation: $\Delta H = 1500 - 1700$.
84. If $\Delta H = -40$ kJ and $\Delta S = -100$ J K$^{-1}$, the process is spontaneous at 200 K because $\Delta G$ is:
Explanation: Convert entropy term: $T\Delta S = 200\times(-0.1) = -20$ kJ, so $\Delta G = -40 - (-20)$.
85. For the same process at 500 K, $\Delta G$ becomes:
Explanation: $\Delta G=-40 - 500(-0.1)=+10$ kJ, so it is not spontaneous.
86. If $\Delta H = +20$ kJ and $\Delta S = +100$ J K$^{-1}$, spontaneity begins above temperature:
Explanation: Set $\Delta G = 0$: $T=\Delta H/\Delta S = 20/0.1$.
87. If 500 g water is heated by 10 K and specific heat is 4.18 J g$^{-1}$ K$^{-1}$, heat absorbed is:
Explanation: $q=ms\Delta T$.
88. If a reaction mixture loses heat to surroundings, the sign of q for the system is:
Explanation: Heat is leaving the system.
89. If surroundings lose heat to the system, q for the system is:
Explanation: Heat enters the system.
90. For the same initial and final states, which can vary with path?
Explanation: Heat and work are path functions.
91. For the same initial and final states, which remains fixed?
Explanation: Internal-energy change depends only on states.
92. In an adiabatic expansion of an ideal gas, temperature usually:
Explanation: The gas does work without absorbing heat.
93. For ideal gas isothermal expansion, change in internal energy is:
Explanation: Internal energy of ideal gas depends only on temperature.
94. For ideal gas, $\Delta H$ depends only on:
Explanation: Enthalpy of an ideal gas is a function of temperature only.
95. Expansion against vacuum produces work equal to:
Explanation: External pressure is zero, so no PV work is done.
96. At the same initial and final states, reversible expansion gives work magnitude:
Explanation: Reversible expansion extracts maximum work.
97. For gaseous reactions, $\Delta H - \Delta U = \Delta n_g RT$ links enthalpy and:
Explanation: It relates the two through gaseous mole change.
98. If $\Delta n_g$ is positive, then for ideal gases:
Explanation: Because $\Delta H = \Delta U + \Delta n_gRT$.
99. Calorimetry problems rely chiefly on the principle of:
Explanation: Heat lost by one part equals heat gained by another.
100. Numerical thermodynamics becomes easy when sign convention, process condition, and the relevant state function are identified before:
Explanation: That prevents the most common mistakes.
101. A spontaneous process at constant T and P must have:
Explanation: That is the Gibbs criterion.
102. At equilibrium, the forward and backward tendencies balance, so:
Explanation: This is the free-energy condition at equilibrium.
103. Among solid, liquid, and gas, the entropy is highest for:
Explanation: Gas particles have maximum freedom of motion.
104. Freezing of water has entropy change for the system:
Explanation: Liquid becomes a more ordered solid.
105. A reaction with negative enthalpy and positive entropy is thermodynamically:
Explanation: Both contributions push $\Delta G$ negative.
106. A reaction with positive enthalpy and positive entropy becomes favorable when:
Explanation: The $T\Delta S$ term can overcome positive $\Delta H$.
107. A reaction with negative enthalpy and negative entropy is favored when:
Explanation: Low T minimizes the unfavorable entropy term.
108. Why is $q_p$ more common in chemistry labs than $q_v$?
Explanation: Laboratory conditions often approximate constant atmospheric pressure.
109. For ideal gases, molar heat capacity at constant pressure is:
Explanation: Extra heat is needed to account for expansion work.
110. For an ideal gas, the difference $C_p - C_v$ equals:
Explanation: This is a standard ideal-gas relation.
111. A reversible process is an idealization in which the system and surroundings differ by:
Explanation: That allows the process to be reversed by tiny changes.
112. Free expansion of a gas into vacuum is:
Explanation: It cannot be retraced by infinitesimal changes.
113. Which is a state function?
Explanation: It depends only on the state of the system.
114. Which is not a state function?
Explanation: Heat depends on how the process occurs.
115. The standard state of an element is the most stable form at:
Explanation: Usually 298 K is implied unless otherwise stated.
116. A fuel with more negative enthalpy of combustion releases:
Explanation: Greater negativity means larger exothermicity.
117. If direct reaction enthalpy is hard to measure, Hess's law suggests using:
Explanation: State-function logic makes the sum valid.
118. The term 'free energy' is used because Gibbs energy measures the portion of energy available to do:
Explanation: More precisely, it corresponds to maximum non-expansion work.
119. A reaction can be exothermic but non-spontaneous at some temperature if:
Explanation: A large unfavorable entropy term can offset negative enthalpy.
120. A reaction can be endothermic yet spontaneous if:
Explanation: The entropy term can drive $\Delta G$ negative.
121. A process with negative $\Delta G$ at constant T and P corresponds to:
Explanation: The two criteria are equivalent under these conditions.
122. For a cyclic process, which pair is possible?
Explanation: State returns to start, though heat and work may be exchanged.
123. A catalyst affects reaction rate but does not change:
Explanation: Catalysts alter kinetics, not thermodynamic state functions.
124. The most frequent sign-convention mistake in thermodynamics is forgetting that work done by the system is:
Explanation: That sign choice is central to correct numerical answers.
125. This chapter is easiest when you sort every question into one of four buckets: energy conservation, heat/enthalpy calculation, entropy, or:
Explanation: That framework covers the whole NEET thermodynamics unit.