Chapter 5: Thermodynamics Short Questions | chemistry Class 11 CBSE

ByteNotes

Chapter 5: Thermodynamics Short Questions | chemistry Class 11 CBSE | ByteNotes.in

Short Answer

Easy difficulty • Direct answer format

Easy

Question 1

Quick Recall

Distinguish between a closed system and an isolated system with one example each.

A closed system allows energy exchange but not matter exchange with the surroundings; example: reactants in a sealed copper or steel vessel.
An isolated system allows neither energy nor matter exchange with the surroundings; example: reactants in a thermos flask or insulated vessel.
The key difference is that a closed system has thermally conducting walls (energy can pass), while an isolated system has insulating walls that prevent all exchange.

ByteNotes

Short Answer

Easy difficulty • Direct answer format

Easy

Question 2

Quick Recall

State the first law of thermodynamics and write its mathematical expression.

The first law of thermodynamics states that the energy of an isolated system is constant, also expressed as the law of conservation of energy: energy can neither be created nor destroyed.
Mathematically: ΔU = q + w, where ΔU is the change in internal energy, q is the heat absorbed by the system, and w is the work done on the system.
By IUPAC convention, q is positive when heat is absorbed by the system and w is positive when work is done on the system.

ByteNotes

Short Answer

Easy difficulty • Direct answer format

Easy

Question 3

Quick Recall

What is a state function? Give two examples and explain why internal energy is a state function.

A state function is a property whose value depends only on the current state of the system and not on the path taken to reach that state.
Examples: internal energy (U), enthalpy (H), temperature (T), volume (V), pressure (p), and entropy (S).
Internal energy is a state function because Joule's experiments showed that the same change in temperature (and hence internal energy) results regardless of whether work is done mechanically or electrically — the path does not matter.

ByteNotes

Short Answer

Easy difficulty • Direct answer format

Easy

Question 4

Quick Recall

What is an adiabatic process? Write the expression for ΔU in an adiabatic process.

An adiabatic process is one in which there is no transfer of heat between the system and the surroundings (q = 0); the system is enclosed by an adiabatic wall.
For an adiabatic process, from the first law: ΔU = q + w = 0 + w = wad, where wad is the adiabatic work done on the system.
This means that any change in internal energy during an adiabatic process is entirely due to work done on or by the system.

ByteNotes

Short Answer

Medium difficulty • Direct answer format

Medium

Question 5

Quick Recall

Define enthalpy. Why is enthalpy preferred over internal energy for reactions carried out at constant pressure?

Enthalpy (H) is a thermodynamic state function defined as H = U + pV; it represents the total heat content of a system.
Most chemical reactions are carried out in open vessels at constant atmospheric pressure, not at constant volume. The heat absorbed at constant pressure equals ΔH (not ΔU).
Since ΔH = qp directly measures the heat exchanged at constant pressure, it is more practical and directly measurable for laboratory reactions than ΔU.

ByteNotes

Short Answer

Medium difficulty • Direct answer format

Medium

Question 6

Quick Recall

Derive the relationship between ΔH and ΔU for reactions involving gases.

At constant pressure: ΔH = ΔU + pΔV; for gaseous reactions, pΔV = ΔngRT, where Δng = moles of gaseous products minus moles of gaseous reactants.
Substituting: ΔH = ΔU + ΔngRT, which is the working equation to interconvert ΔH and ΔU.
When Δng = 0, ΔH = ΔU; when Δng > 0, ΔH > ΔU; when Δng < 0, ΔH < ΔU.

ByteNotes

Short Answer

Medium difficulty • Direct answer format

Medium

Question 7

Quick Recall

What is a bomb calorimeter? What thermodynamic quantity does it measure?

A bomb calorimeter is a sealed steel vessel immersed in a water bath, used to measure heat changes at constant volume by burning a combustible substance in pure dioxygen.
Since the volume is constant (ΔV = 0), no pressure-volume work is done, so the heat measured directly gives qV = ΔU, the change in internal energy.
The temperature change of the water surrounding the bomb is monitored, and using the known heat capacity of the calorimeter, qV is calculated using q = CΔT.

ByteNotes

Short Answer

Medium difficulty • Direct answer format

Medium

Question 8

Quick Recall

Define standard enthalpy of formation. Why is the standard enthalpy of formation of an element in its reference state taken as zero?

Standard enthalpy of formation (ΔfH°) is the enthalpy change when one mole of a compound is formed from its constituent elements in their most stable reference states at 1 bar and specified temperature (usually 298 K).
The reference state of an element is its most stable form at 25°C and 1 bar (e.g., H2 gas, O2 gas, graphite for carbon, rhombic sulphur).
Since no compound is being formed in this case (element → same element), there is no chemical change, so by convention ΔfH° = 0 for elements in their reference states.

ByteNotes

Short Answer

Medium difficulty • Direct answer format

Medium

Question 9

Quick Recall

State Hess's law of constant heat summation and give its thermodynamic basis.

Hess's law states that if a reaction takes place in several steps, its standard reaction enthalpy is the sum of the standard enthalpies of the intermediate reactions at the same temperature.
The thermodynamic basis is that enthalpy is a state function — the total enthalpy change depends only on the initial and final states, not on the path taken.
Mathematically: ΔrH = ΔrH1 + ΔrH2 + ΔrH3 + ..., allowing calculation of enthalpy changes for reactions that cannot be measured directly (like partial combustion).

ByteNotes

Short Answer

Medium difficulty • Direct answer format

Medium

Question 10

Quick Recall

Differentiate between bond dissociation enthalpy and mean bond enthalpy with reference to methane.

Bond dissociation enthalpy is the enthalpy change for breaking a specific bond in a molecule; for diatomic molecules, it equals the enthalpy of atomization.
Mean bond enthalpy is the average of bond dissociation enthalpies for the same type of bond broken successively in a polyatomic molecule.
In CH4, the four C-H bonds have different individual dissociation enthalpies (+427, +439, +452, +347 kJ mol-1 respectively), so mean C-H bond enthalpy is calculated as ΔaH/4 = 1665/4 = 416 kJ mol-1.

ByteNotes

Short Answer

Medium difficulty • Direct answer format

Medium

Question 11

Quick Recall

What is entropy? State how entropy changes are qualitatively predicted for physical state changes.

Entropy (S) is a thermodynamic state function that measures the degree of disorder or randomness in a system; greater disorder corresponds to higher entropy.
For a given substance: the crystalline solid state has the lowest entropy (most ordered), the liquid state has intermediate entropy, and the gaseous state has the highest entropy.
Phase changes that increase disorder (melting, vaporization, sublimation) involve a positive ΔS, while processes that increase order (crystallization, condensation) involve a negative ΔS.

ByteNotes

Short Answer

Medium difficulty • Direct answer format

Medium

Question 12

Quick Recall

State the second law of thermodynamics and express the condition for a spontaneous process in terms of total entropy.

The second law of thermodynamics states that the entropy of an isolated system always tends to increase, and the natural direction of spontaneous change is toward greater disorder.
For a spontaneous process, the total entropy change (ΔStotal = ΔSsystem + ΔSsurroundings) must be greater than zero: ΔStotal > 0.
At equilibrium, ΔStotal = 0; for non-spontaneous (reverse) processes, ΔStotal < 0. The second law explains why spontaneous exothermic reactions are common — they increase surroundings entropy.

ByteNotes

Short Answer

Medium difficulty • Direct answer format

Medium

Question 13

Quick Recall

Define Gibbs energy and write the expression for its change at constant temperature.

Gibbs energy (G) is a thermodynamic state function defined as G = H - TS, combining enthalpy and entropy of a system; it represents the net energy available to do useful work.
At constant temperature, the change in Gibbs energy is given by: ΔG = ΔH - TΔS, known as the Gibbs equation.
ΔG < 0 indicates spontaneous process; ΔG = 0 indicates equilibrium; ΔG > 0 indicates non-spontaneous process.

ByteNotes

Premium Access

Unlock all Short Questions

You are viewing 13 free short questions from this chapter. Upgrade to Premium to unlock the remaining 12 questions.

Upgrade to Premium12 questions locked

Chapter sections

Chapter 5: Thermodynamics | Short Questions

Distinguish between a closed system and an isolated system with one example each.

State the first law of thermodynamics and write its mathematical expression.

What is a state function? Give two examples and explain why internal energy is a state function.

What is an adiabatic process? Write the expression for ΔU in an adiabatic process.

Define enthalpy. Why is enthalpy preferred over internal energy for reactions carried out at constant pressure?

Derive the relationship between ΔH and ΔU for reactions involving gases.

What is a bomb calorimeter? What thermodynamic quantity does it measure?

Define standard enthalpy of formation. Why is the standard enthalpy of formation of an element in its reference state taken as zero?

State Hess's law of constant heat summation and give its thermodynamic basis.

Differentiate between bond dissociation enthalpy and mean bond enthalpy with reference to methane.

What is entropy? State how entropy changes are qualitatively predicted for physical state changes.

State the second law of thermodynamics and express the condition for a spontaneous process in terms of total entropy.

Define Gibbs energy and write the expression for its change at constant temperature.

Unlock all Short Questions