78 free MCQs on Thermodynamics, each with its own worked answer and explanation. Study energy changes in chemical reactions. Understand enthalpy, entropy, Gibbs free energy, and the laws of thermodynamics that decide whether a reaction will occur spontaneously or not.
Take the timed Thermodynamics chapterwise test →78 practice questions on Thermodynamics, sorted Easy → Hard. Try each one first, then open its answer page for the worked explanation. Want the full theory first? Read the Thermodynamics notes.
Isothermal expansion follows a gentler curve (heat flows in to keep T constant) while adiabatic expansion drops in pressure more steeply (no heat exchange, so internal energy and temperature fall as the gas does work).
Q1.
Which of the following is an intensive property?
Q2.
According to the equation ΔU = q + w, when work is done on the system, the sign of w is:
Q3.
Which of the following is a path function?
Q4.
In NCERT thermodynamics, the standard state of a substance is defined at a pressure of:
Q5.
For a complete cyclic process, the change in internal energy (ΔU) is:
Q6.
A reaction that releases heat to its surroundings is called:
Q8.
Which law of thermodynamics states that energy cannot be created or destroyed?
Q9.
Enthalpy change at constant pressure equals:
Q10.
A reaction is spontaneous when Gibbs free energy change (ΔG) is:
Q11.
The standard enthalpy of formation of any element in its most stable state is:
Q12.
In an adiabatic process:
Q14.
Gibbs free energy is defined as:
Q15.
Which thermodynamic process occurs at constant temperature?
Q16.
The branch of science dealing with energy changes accompanying chemical and physical processes is:
Q17.
A system that exchanges both energy and matter with its surroundings is a(n):
Q18.
An ideal thermos flask, exchanging neither matter nor energy, is a(n):
Q19.
Which of the following is a state function?
Q20.
The first law of thermodynamics is a statement of the conservation of:
Q21.
For an exothermic reaction, the enthalpy change ΔH is:
Q23.
At constant pressure, the heat exchanged by a system equals its change in:
Q24.
The internal energy of a system is represented by the symbol:
Q25.
A process carried out at constant temperature is described as:
Q26.
In a bomb calorimeter the reaction is carried out at constant volume, so the heat measured gives directly:
Q27.
For H<sub>2</sub>(g) + Cl<sub>2</sub>(g) → 2HCl(g), using bond enthalpies H–H = 435, Cl–Cl = 242 and H–Cl = 431 kJ/mol, ΔH is:
Q28.
A system absorbs 700 J of heat and does 300 J of work on the surroundings. Its change in internal energy is:
Q29.
Given ΔHf(CO<sub>2</sub>) = −393.5 and ΔHf(CO) = −110.5 kJ/mol, the enthalpy change for CO(g) + ½O<sub>2</sub>(g) → CO<sub>2</sub>(g) is:
Q30.
The enthalpy of vaporisation of water is 40.8 kJ/mol at its boiling point of 373 K. The entropy of vaporisation is about:
Q31.
For which of these reactions is ΔH equal to ΔU (i.e. Δn<sub>g</sub> = 0)?
Q32.
For a reaction with ΔH = -200 kJ and ΔS = +100 J/K at 300 K, ΔG is:
Q33.
In an isothermal reversible expansion of an ideal gas, ΔU is:
Q34.
Bond enthalpy (bond dissociation energy) is always:
Q35.
Hess's law states that the enthalpy change of a reaction:
Q36.
The entropy change for melting of ice at 0°C (ΔH_fusion = 6 kJ/mol) is:
Q37.
A reaction is spontaneous only at high temperatures when:
Q38.
The standard Gibbs energy change relates to the equilibrium constant by:
Q39.
Work done during isothermal reversible expansion of an ideal gas is:
Q40.
For an ideal gas undergoing free expansion into a vacuum, the values of q and w are:
Q41.
For the reaction N<sub>2</sub>(g) + 3H<sub>2</sub>(g) -> 2NH<sub>3</sub>(g), how is the standard enthalpy of formation of NH<sub>3</sub> related to the standard reaction enthalpy?
Q42.
For the isothermal expansion of an ideal gas, the change in internal energy ΔU is:
Q43.
In an adiabatic process, the heat exchanged with the surroundings (q) is:
Q44.
In ΔH = ΔU + ΔnₘRT, the term Δnₘ is the change in the number of moles of:
Q45.
Which of these has the highest standard molar entropy at 298 K?
Q46.
Hess’s law of constant heat summation is a direct consequence of enthalpy being a:
Q47.
The enthalpy change when one mole of a compound forms from its elements in their standard states is the standard enthalpy of:
Q48.
For the reaction N₂(g) + 3H₂(g) → 2NH₃(g), the value of Δnₘ (gaseous) is:
Q49.
The heat change measured at constant volume is equal to the change in:
Q50.
For one mole of an ideal gas, the difference between the molar heat capacities Cp and Cv equals:
Q51.
For the isothermal reversible expansion of an ideal gas from V₁ to V₂, the work done by the gas is:
Q52.
1 g of graphite is burnt in a bomb calorimeter of heat capacity 20.7 kJ K<sup>−1</sup> and the temperature rises from 298 K to 299 K. ΔU per mole of carbon is:
Q53.
Given ΔHf(CO<sub>2</sub>) = −393.5, ΔHf(H<sub>2</sub>O, l) = −285.8 and ΔHf(C<sub>2</sub>H<sub>5</sub>OH, l) = −277.0 kJ/mol, the standard enthalpy of combustion of ethanol is:
Q54.
A gas expands from 2 L to 6 L against a constant external pressure of 2 atm. The work done by the gas is about (1 L·atm = 101.3 J):
Q55.
For a reaction with Δn<sub>g</sub> = +2 carried out at 300 K, the value of (ΔH − ΔU) is (R = 8.314 J/K/mol):
Q56.
Assertion (A): In the isothermal reversible expansion of an ideal gas, q = −w. Reason (R): The internal energy of an ideal gas depends only on temperature, so ΔU = 0 in an isothermal change.
Q57.
An exothermic reaction releases 100 kJ of heat to the surroundings at 300 K. The entropy change of the surroundings is approximately:
Q58.
At what temperature does a reaction with ΔH = +80 kJ/mol and ΔS = +200 J/K/mol become spontaneous?
Q59.
For C(s) + O₂(g) → CO₂(g), the relationship between ΔH and ΔU is:
Q60.
Given ΔHc(C) = -393.5, ΔHc(H₂) = -285.8, ΔHc(CH₄) = -890 kJ/mol, the enthalpy of formation of CH₄ is:
Q61.
For a reaction at 298 K with ΔG° = -20 kJ/mol, the equilibrium constant K is approximately:
Q62.
The Kirchhoff equation d(ΔH)/dT = ΔCp is used to calculate:
Q63.
For an adiabatic reversible expansion, which relation holds?
Q64.
The efficiency of a Carnot engine operating between 500 K and 300 K is:
Q65.
Which of the following has a negative entropy change (ΔS < 0)?
Q66.
For the reaction N<sub>2</sub>(g) + 3H<sub>2</sub>(g) -> 2NH<sub>3</sub>(g) at 300 K, if the internal energy change (Delta-U) is -100 kJ, what is the enthalpy change Delta-H (R = 8.314 J/K/mol)?
Q67.
One mole of an ideal gas is compressed isothermally and reversibly from 10 L to 1 L at 300 K. What is the work done on the gas (R = 8.314 J/K/mol)?
Q68.
The second law of thermodynamics states that for any spontaneous process the total entropy of the universe:
Q69.
The Gibbs free energy change is defined by the equation ΔG =:
Q70.
A reaction is spontaneous at all temperatures when:
Q71.
At equilibrium, the Gibbs free energy change ΔG for the reaction is:
Q72.
The relationship between the standard free energy change and the equilibrium constant is ΔG° =:
Q73.
A reaction has ΔH = +50 kJ/mol and ΔS = +100 J/K·mol. It becomes spontaneous above a temperature of about:
Q74.
The entropy change for the melting of ice at 0 °C is:
Q75.
By the third law of thermodynamics, the entropy of a perfect crystal at absolute zero is:
Q76.
For a reversible adiabatic process involving an ideal gas, which relation holds?
Q77.
A Carnot engine operating between 500 K and 300 K has an efficiency of:
Q78.
A reaction in a bomb calorimeter gives ΔU = −742.7 kJ mol<sup>−1</sup> at 298 K with Δn<sub>g</sub> = −1. Taking R = 8.314 J K<sup>−1</sup> mol<sup>−1</sup>, ΔH is:
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