Are these practice questions for Thermodynamics based on the CBSE pattern?

Yes. Every practice question is modeled on the CBSE Class 11 - Science Chemistry exam pattern — including MCQs, assertion-reason questions, and short-answer formats that frequently appear in CBSE papers.

Are the practice questions for Thermodynamics free?

Yes. Every chapter includes free practice questions for Class 11 - Science students to get you started.

What types of questions are included?

The practice set for Thermodynamics mainly includes MCQs and CBSE-style short-answer questions, with instant feedback on each answer. For full-length textbook practice, use the NCERT Solutions page.

Are the practice questions suitable for board exam preparation?

Yes. The questions test the same concepts and depth that CBSE asks in board exams. After finishing, also solve the NCERT and NCERT Exemplar exercises for Thermodynamics for full coverage.

How many practice questions are there for Thermodynamics?

The number scales with chapter length and concept density. Each chapter typically has 15-50 practice questions across MCQ and short-answer formats.

Thermodynamics Class 11 Chemistry Practice Questions with Answers

easy

Explain the difference between a system and its surroundings in thermodynamics, using an example.

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easy

For the reaction

2SO_2(g) + O_2(g) \rightarrow 2SO_3(g)

, analyze whether

\Delta H

is greater than, less than, or equal to

\Delta U

.

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easy

Define an isolated system.

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easy

State the first law of thermodynamics.

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easy

Define entropy (S).

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easy

What is an intensive property? Give one example.

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easy

Propose a reason why the standard enthalpy of formation of C(diamond) is not zero, while that of C(graphite) is zero.

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easy

A system absorbs 500 J of heat and does 200 J of work on the surroundings. Calculate the change in internal energy of the system.

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easy

What is the value of the standard enthalpy of formation for an element in its reference state?

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easy

Calculate the work done when 2 moles of an ideal gas expand from a volume of 5 L to 15 L against a constant external pressure of 2 atm. (1 L atm = 101.3 J)

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easy

Calculate the amount of heat required to raise the temperature of 100 g of copper from

25^\circ\text{C}

to

125^\circ\text{C}

. The specific heat capacity of copper is

0.385 \text{ J g}^{-1} \text{ K}^{-1}

.

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medium

Formulate a Born-Haber cycle for the formation of solid magnesium oxide (MgO). From this cycle, create a single algebraic expression to represent the lattice enthalpy (

\Delta_{\text{lattice}}H^\ominus

) of MgO using the following defined thermodynamic terms:

Enthalpy of formation of MgO(s): $\Delta_f H^\ominus$
Enthalpy of sublimation of Mg(s): $\Delta_{\text{sub}}H^\ominus$
First and second ionization enthalpies of Mg(g): $IE_1$ and $IE_2$
Bond dissociation enthalpy of O₂(g): $\Delta_{\text{bond}}H^\ominus$
First and second electron gain enthalpies of O(g): $\Delta_{eg}H_1^\ominus$ and $\Delta_{eg}H_2^\ominus$

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medium

Explain the sign conventions for heat (q) and work (w) in chemical thermodynamics as per IUPAC.

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medium

For the reaction

N_2O_4(g) \rightleftharpoons 2NO_2(g)

, the standard Gibbs energy change

\Delta_r G^\ominus

is

+4.85 \text{ kJ mol}^{-1}

at

298 \text{ K}

. (a) Calculate the equilibrium constant,

K_p

. (b) Analyze whether the formation of

NO_2

is favored under standard conditions. (

R = 8.314 \text{ J K}^{-1} \text{ mol}^{-1}

)

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medium

A student claims that for the reaction

2\text{SO}_2(\text{g}) + \text{O}_2(\text{g}) \rightarrow 2\text{SO}_3(\text{g})

, the values of

\Delta H

and

\Delta U

are nearly identical. Critique this reasoning. Then, assuming the reaction occurs at 298 K and has

\Delta_r H^\ominus = -198.0 \text{ kJ mol}^{-1}

, calculate the value of

\Delta_r U^\ominus

and justify whether the student's claim is valid.

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medium

Explain the following thermodynamic terms: (a) Adiabatic process (b) Extensive properties (c) Reversible process (d) Spontaneous process

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medium

State Hess's Law of Constant Heat Summation and explain its importance.

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medium

Evaluate the spontaneity of the vaporization of a liquid at its normal boiling point. Justify your conclusion by analyzing the signs of

\Delta H

,

\Delta S

, and the resulting value of

\Delta G

.

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medium

Examine the following scenarios and classify each system as open, closed, or isolated: (a) Coffee in a perfect thermos flask. (b) A human being.

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medium

Differentiate between open, closed, and isolated systems. Provide one example for each type.

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medium

(a) Define standard enthalpy of formation (

\Delta_f H^{\ominus}

). (b) Explain what a thermochemical equation is, providing a complete example.

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medium

Propose a complete experimental and computational procedure to determine the standard enthalpy of formation of liquid benzene (

\text{C}_6\text{H}_6(\text{l})

) using only data obtainable from bomb calorimetry. You are provided with pure samples of benzene, high-purity graphite, and hydrogen gas. You do not need to provide numerical values, but you must formulate the specific thermochemical equations and the final Hess's Law calculation.

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medium

Calculate the change in internal energy (

\Delta U

) when 1 mole of liquid water vaporizes at

100^\circ\text{C}

and 1 atm pressure. The enthalpy of vaporization (

\Delta_{vap}H

) is

40.79 \text{ kJ mol}^{-1}

. Assume the water vapor behaves as an ideal gas. (

R = 8.314 \text{ J K}^{-1} \text{ mol}^{-1}

)

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medium

Justify why the enthalpy of neutralization for a strong acid reacting with a strong base is a constant value, whereas it varies when a weak acid or weak base is involved.

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medium

2 moles of an ideal gas expand isothermally and reversibly from 1 L to 10 L at

27^\circ\text{C}

. Calculate the work done by the gas. (

R = 8.314 \text{ J K}^{-1} \text{ mol}^{-1}

)

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medium

Critique the statement: "All exothermic reactions are spontaneous." Use thermodynamic principles to justify your conclusion.

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medium

Formulate the general relationship between molar heat capacity at constant pressure (

C_p

) and molar heat capacity at constant volume (

C_v

) for an ideal gas. Starting from the fundamental definitions of enthalpy and heat capacity, derive the expression

C_p - C_v = R

. Justify each mathematical step in your derivation.

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medium

Justify, using the concept of pressure-volume work, why the magnitude of work done in a reversible isothermal expansion of an ideal gas is greater than that in an irreversible isothermal expansion between the same initial and final volumes.

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medium

The combustion of 1.2 g of benzoic acid (

C_7H_6O_2

) in a bomb calorimeter caused a temperature rise of

2.15 \text{ K}

. The heat capacity of the calorimeter is

14.8 \text{ kJ K}^{-1}

. Calculate the molar enthalpy of combustion (

\Delta_c H^\ominus

) of benzoic acid at

298 \text{ K}

.

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medium

The enthalpy of fusion for ice is

6.01 \text{ kJ mol}^{-1}

. Calculate the heat required to melt 9 g of ice.

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medium

Define the terms: (a) Internal Energy (U), (b) Enthalpy (H), and (c) State Function.

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medium

Compare the magnitude of work done by the system during the expansion of a gas from

V_1

to

V_2

under (a) isothermal reversible conditions and (b) isothermal irreversible conditions (against a constant external pressure,

p_{ex}

). Demonstrate which process yields more work.

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medium

For a reaction,

\Delta H = -28.0 \text{ kJ mol}^{-1}

and

\Delta S = -60.0 \text{ J K}^{-1} \text{ mol}^{-1}

. At what temperature will the reaction become spontaneous? Analyze the conditions for spontaneity.

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medium

What is meant by the standard enthalpy of vaporization? Explain why the enthalpy of vaporization for water is higher than that for acetone.

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hard

Evaluate whether the Gibbs equation,

\Delta G = \Delta H - T\Delta S

, can be applied to a process that occurs under non-isothermal conditions. Justify your answer.

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hard

An engineer is designing a process for the synthesis of ammonia:

\text{N}_2(\text{g}) + 3\text{H}_2(\text{g}) \rightleftharpoons 2\text{NH}_3(\text{g})

. The standard enthalpy of formation (

\Delta_f H^\ominus

) for

\text{NH}_3(\text{g})

is

-46.1 \text{ kJ mol}^{-1}

and the standard entropy change for the reaction (

\Delta_r S^\ominus

) is

-198.7 \text{ J K}^{-1}\text{mol}^{-1}

. (a) Create an expression for the standard Gibbs free energy change (

\Delta_r G^\ominus

) as a function of temperature (T). (b) Evaluate the temperature at which the reaction is at equilibrium under standard conditions. (c) Justify, using thermodynamic principles, why the industrial Haber-Bosch process operates at moderately high temperatures (e.g., 400-450°C) and high pressures, despite low temperatures favoring spontaneity.

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hard

Critique the statement: "A reaction with a positive

\Delta_r H^\ominus

and a negative

\Delta_r S^\ominus

can never be spontaneous." Justify your answer using the Gibbs free energy equation. Then, consider the reverse reaction. Formulate the expressions for its

\Delta_r H^\ominus

,

\Delta_r S^\ominus

, and

\Delta_r G^\ominus

, and evaluate its spontaneity.

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hard

Justify that the free expansion of an ideal gas is a spontaneous process by analyzing the entropy changes involved. In your justification, evaluate

\Delta U, q, w

, and

\Delta S_{\text{system}}

,

\Delta S_{\text{surroundings}}

, and

\Delta S_{\text{total}}

for this irreversible isothermal process.

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hard

Calculate the standard enthalpy of formation (

\Delta_f H^\ominus

) of liquid methanol (

CH_3OH(l)

) from the following data:

$\Delta_f H^\ominus (CO_2, g) = -393.5 \text{ kJ mol}^{-1}$
$\Delta_f H^\ominus (H_2O, l) = -285.8 \text{ kJ mol}^{-1}$
The standard enthalpy of combustion of methanol is $\Delta_c H^\ominus = -726.0 \text{ kJ mol}^{-1}$ . First, write the balanced thermochemical equation for the combustion of methanol. Then, use the data to solve for the enthalpy of formation.

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hard

Explain the concept of Gibbs energy (G) and its role as a criterion for spontaneity. (a) Write the Gibbs equation. (b) Describe the conditions for spontaneity and equilibrium in terms of

\Delta G

. (c) Explain why an endothermic process (positive

\Delta H

) can still be spontaneous.

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hard

Using the following thermochemical equations, calculate the standard enthalpy of formation (

\Delta_f H^\ominus

) for acetylene,

C_2H_2(g)

.

$C(\text{graphite}) + O_2(g) \rightarrow CO_2(g)$ ; $\Delta H^\ominus = -393.5 \text{ kJ mol}^{-1}$
$H_2(g) + \frac{1}{2}O_2(g) \rightarrow H_2O(l)$ ; $\Delta H^\ominus = -285.8 \text{ kJ mol}^{-1}$
$C_2H_2(g) + \frac{5}{2}O_2(g) \rightarrow 2CO_2(g) + H_2O(l)$ ; $\Delta H^\ominus = -1299.6 \text{ kJ mol}^{-1}$

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hard

(a) State the mathematical relationship between enthalpy change (

\Delta H

) and internal energy change (

\Delta U

). (b) Derive the expression

\Delta H = \Delta U + \Delta n_g RT

. (c) Under what condition is

\Delta H

equal to

\Delta U

?

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hard

Design a hypothetical two-step process to calculate the standard enthalpy of formation of carbon monoxide (CO) using Hess's Law. You are given the standard enthalpies of combustion for carbon (graphite) and carbon monoxide. Justify why a direct experimental measurement of the enthalpy of formation of CO is difficult.

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hard

Given the reaction:

\text{N}_2\text{O}_4(\text{g}) \rightleftharpoons 2\text{NO}_2(\text{g})

. At 350 K, the equilibrium constant

K_p

is 2.5. A chemist proposes that increasing the temperature will favor the formation of

\text{NO}_2

. Evaluate this proposal, given that the standard enthalpy of reaction,

\Delta_r H^\ominus

, is

+57.2 \text{ kJ mol}^{-1}

. Justify your answer using the relationship between Gibbs free energy, enthalpy, and the equilibrium constant.

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hard

Calculate the standard enthalpy of reaction (

\Delta_r H^\ominus

) for the hydrogenation of ethene:

C_2H_4(g) + H_2(g) \rightarrow C_2H_6(g)

. Use the following bond enthalpy data:

$C-H$ : $414 \text{ kJ mol}^{-1}$
$C=C$ : $611 \text{ kJ mol}^{-1}$
$H-H$ : $436 \text{ kJ mol}^{-1}$
$C-C$ : $347 \text{ kJ mol}^{-1}$

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ThermodynamicsClass 11 Chemistry Practice Questions