Complete List of Theoretical Questions

Section A: Basic Thermodynamic Concepts

• What is thermodynamics?
• What is a thermodynamic system?
• What are surroundings?
• Define the boundary of a system.
• Differentiate between system and surroundings.
• What is meant by the universe in thermodynamics?
• Differentiate between open, closed and isolated systems with examples.
• What is the thermodynamic state of a system?
• What are state variables or state functions?
• Why are pressure, volume and temperature called state functions?
• What is meant by equilibrium state?
• Why does thermodynamics deal with macroscopic properties?
• What is an adiabatic system?
• What is an adiabatic process?
• What is an adiabatic wall?
• Differentiate between adiabatic and isothermal processes.

Section B: Internal Energy, Heat and Work

• Define internal energy.
• What are the different forms of energy contributing to internal energy?
• Why is internal energy called a state function?
• Why can we measure only the change in internal energy and not its absolute value?
• Define heat in thermodynamics.
• Define work in thermodynamics.
• Differentiate between heat and work.
• Explain the sign conventions of heat and work according to IUPAC.
• When is heat considered positive and negative?
• When is work considered positive and negative?
• State the first law of thermodynamics.
• Give the mathematical expression of the first law of thermodynamics.
• Explain the law of conservation of energy.
• Why are heat and work called path functions?
• Why are heat and work not state functions?
• Can the internal energy of an isolated system change?
• Explain the relationship between heat, work and internal energy.
• What is pressure-volume work?
• What is free expansion of a gas?
• Why is no work done during free expansion into vacuum?
• What is reversible expansion?
• What is irreversible expansion?
• Differentiate between reversible and irreversible processes.
• Why is a reversible process infinitely slow?
• Why is the work done in reversible expansion greater than in irreversible expansion under the same initial and final conditions?

Section C: Enthalpy

• What is enthalpy?
• Why do we introduce enthalpy as a new thermodynamic state function?
• Derive the relationship between enthalpy and internal energy.
• Why is enthalpy a state function?
• What is the physical significance of enthalpy change?
• What is the relationship between heat absorbed at constant pressure and enthalpy change?
• What is the relationship between heat absorbed at constant volume and internal energy change?
• Differentiate between ΔH and ΔU.
• Establish the relationship:

\Delta H=\Delta U+\Delta n_gRT

• Under what conditions is ΔH equal to ΔU?
• Why is the difference between ΔH and ΔU generally insignificant for solids and liquids?
• Why does the difference between ΔH and ΔU become significant when gases are involved?
• What is an exothermic reaction?
• What is an endothermic reaction?
• Explain the sign of enthalpy change for exothermic and endothermic reactions.

Section D: Extensive and Intensive Properties

• What are extensive properties?
• What are intensive properties?
• Differentiate between extensive and intensive properties with examples.
• Why is internal energy an extensive property?
• Why are temperature and pressure intensive properties?
• What is a molar property?
• Why is molar volume an intensive property?

Section E: Heat Capacity

• What is heat capacity?
• Define specific heat capacity.
• Define molar heat capacity.
• Differentiate between heat capacity and specific heat capacity.
• What is the relationship between heat, heat capacity and temperature change?
• What is the difference between and ?
• Derive the relationship between and for an ideal gas.

C_p-C_v=R

• Why is greater than for an ideal gas?
• Why does water have a large heat capacity?

Section F: Measurement of Enthalpy and Internal Energy

• What is calorimetry?
• What is a calorimeter?
• Explain the construction and working of a bomb calorimeter.
• Why is a bomb calorimeter used to measure ΔU?
• Why is a bomb calorimeter operated at constant volume?
• How is enthalpy change measured at constant pressure?
• Differentiate between bomb calorimeter and constant-pressure calorimeter.
• Why is the heat lost by the reaction equal to the heat gained by the calorimeter?

Section G: Enthalpy Changes of Reactions

• What is reaction enthalpy?
• What is standard enthalpy of reaction?
• What is the standard state of a substance?
• Why is 1 bar pressure used as the standard pressure?
• What is the standard temperature generally used for thermodynamic data?
• What is standard enthalpy of fusion?
• What is standard enthalpy of vaporisation?
• What is standard enthalpy of sublimation?
• Differentiate between fusion, vaporisation and sublimation.
• Why are enthalpies of fusion and vaporisation positive?
• Why does temperature remain constant during a phase transition?
• What factors affect the enthalpy change during phase transformation?
• Why does water have a higher enthalpy of vaporisation than acetone?
• What is standard enthalpy of formation?
• Why is the standard enthalpy of formation of an element in its reference state taken as zero?
• What is a reference state?
• Differentiate between standard enthalpy of formation and standard enthalpy of reaction.
• What is a thermochemical equation?
• Write the important conventions of thermochemical equations.
• Why must physical states be mentioned in thermochemical equations?

Section H: Different Types of Enthalpy

• What is standard enthalpy of combustion?
• Why are combustion reactions generally exothermic?
• What is enthalpy of atomisation?
• What is bond dissociation enthalpy?
• What is mean bond enthalpy?
• Differentiate between bond dissociation enthalpy and mean bond enthalpy.
• Why is energy required to break a chemical bond?
• Why is energy released during bond formation?
• Why does bond dissociation enthalpy differ for different bonds in a polyatomic molecule?
• What is the relationship between enthalpy of atomisation and bond dissociation enthalpy?
• How can bond enthalpies be used to calculate reaction enthalpy?

Section I: Hess’s Law

• State Hess’s law of constant heat summation.
• Explain Hess’s law with a suitable example.
• Why is Hess’s law based on the state-function nature of enthalpy?
• Why is the enthalpy change independent of the reaction pathway?
• What happens to the enthalpy change when a thermochemical equation is reversed?
• What happens to enthalpy change when the coefficients of a thermochemical equation are multiplied by a number?
• Explain the importance of Hess’s law.
• How can Hess’s law be used to calculate enthalpy changes that cannot be measured experimentally?

Section J: Spontaneity and Entropy

• What is a spontaneous process?
• What is a non-spontaneous process?
• Does a spontaneous process always occur rapidly?
• Differentiate between spontaneous and non-spontaneous processes.
• Why does heat flow spontaneously from a hotter body to a colder body?
• Why does a gas spontaneously expand into an available vacuum?
• Why is the first law of thermodynamics insufficient to predict the spontaneity of a reaction?
• What is the driving force of a spontaneous process?
• What is entropy?
• Why is entropy called a measure of randomness or disorder?
• Why is entropy a state function?
• What is the relationship between entropy change and heat transfer in a reversible process?
• State the second law of thermodynamics.
• What is the entropy change of the universe for a spontaneous process?
• What is the entropy change of an isolated system during a spontaneous process?
• Can a spontaneous process have a negative entropy change of the system?
• Why can an exothermic reaction be non-spontaneous?
• Why can an endothermic reaction be spontaneous?
• Explain entropy change during melting, vaporisation and sublimation.
• Why does entropy increase when a solid changes into a liquid?
• Why does entropy increase when a liquid changes into a gas?
• What is absolute entropy?
• State the third law of thermodynamics.
• What is the entropy of a perfectly crystalline substance at absolute zero?

Section K: Gibbs Energy

• What is Gibbs energy?
• Why is Gibbs energy also called Gibbs free energy?
• Write the mathematical expression for Gibbs energy.
• Derive the Gibbs equation:

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

• Why is Gibbs energy a state function?
• What is the relationship between Gibbs energy change and spontaneity?
• What is the condition for a reaction to be spontaneous?
• What is the condition for a reaction to be non-spontaneous?
• What is the condition for a system at equilibrium?
• Explain the significance of negative, positive and zero values of ΔG.
• Why is Gibbs energy useful for predicting spontaneity at constant temperature and pressure?
• How does temperature affect the spontaneity of a reaction?
• Explain the four possible combinations of ΔH and ΔS for spontaneity.
• Under what conditions can an endothermic reaction become spontaneous?
• Under what conditions can an exothermic reaction become non-spontaneous?

Section L: Gibbs Energy and Equilibrium

• What is standard Gibbs energy change?
• What is the relationship between standard Gibbs energy change and equilibrium constant?
• Derive the relationship:

DeltaG= -RT\ln K

• What is the significance of the magnitude of the equilibrium constant?
• What does a large positive value of ΔG° indicate about the equilibrium constant?
• What does a large negative value of ΔG° indicate about the equilibrium constant?
• What is the relationship between ΔG° and the extent of a chemical reaction?
• How can the equilibrium constant be calculated from standard Gibbs energy change?
• Why does temperature play an important role in determining spontaneity?

Section M: Important NCERT Exercise-Based Theory Questions

• What is the correct condition for a process to occur under adiabatic conditions?
• Why are the enthalpies of elements in their standard states taken as zero for formation enthalpy calculations?
• How does positive entropy change affect the spontaneity of a reaction?
• For an isolated system, if ΔU = 0, what can be said about ΔS?
• What are the signs of ΔH and ΔS for the combination of gaseous chlorine atoms to form chlorine gas?
• Explain the thermodynamic stability of nitric oxide (NO).
• Can a reaction have positive ΔH and still be spontaneous? Explain.
• Can a reaction have negative ΔS and still be spontaneous? Explain.