Sample Quizzes For Preparation: Chemical Energetics

A2 Level Chemistry – Chapter 23: Chemical Energetics Quiz

Question 1: What is lattice energy (ΔHlatt)?

A. Energy required to break one mole of covalent bonds
B. Energy released when gaseous ions form one mole of ionic solid
C. Energy change when a gas becomes a liquid
D. Energy required to form gaseous atoms from an element

Question 2: What is the enthalpy change of atomisation of chlorine?

A. Cl₂(g) → 2Cl⁻(g)
B. ½Cl₂(g) → Cl(g)
C. Cl⁻(g) → Cl(g)
D. Cl₂(l) → Cl₂(g)

Question 3: What does a Born–Haber cycle calculate?

A. Activation energy
B. Gibbs free energy
C. Lattice energy
D. Reaction rate

Question 4: Which factor increases lattice energy (makes it more exothermic)?

A. Larger ionic radius
B. Lower nuclear charge
C. Smaller ionic radius and higher ionic charge
D. Low electronegativity

Question 5: Which of the following is an exothermic process?

A. Ionisation energy
B. Atomisation
C. Lattice formation
D. Breaking covalent bonds

Question 6: What is the standard first electron affinity of chlorine?

A. Cl(g) → Cl⁺(g) + e⁻
B. Cl⁻(g) → Cl(g) + e⁻
C. Cl(g) + e⁻ → Cl⁻(g)
D. Cl₂(g) → 2Cl(g)

Question 7: Why is the second electron affinity of oxygen endothermic?

A. Electron enters a new orbital
B. Energy is needed to overcome repulsion from negative ion
C. O⁻ is stable
D. Electron enters a full sub-shell

Question 8: Which combination gives the most exothermic ΔHlatt?

A. Na⁺ and Cl⁻
B. Mg²⁺ and O²⁻
C. K⁺ and I⁻
D. Ca²⁺ and Br⁻

Question 9: Which process is part of the Born–Haber cycle?

A. Dissolving the salt
B. Neutralisation reaction
C. Atomisation of elements
D. Electrolysis

Question 10: In the cycle, which quantity is typically the most endothermic?

A. Electron affinity
B. Atomisation of metal
C. Ionisation energy
D. Lattice formation

Question 11: What is the enthalpy change of hydration?

A. Energy released when gaseous ions form solid
B. Energy released when gaseous ions become aqueous
C. Energy required to evaporate water
D. Energy change when water boils

Question 12: Which factor increases hydration enthalpy?

A. Larger ionic radius
B. Lower ionic charge
C. Higher charge density
D. Greater molar mass

Question 13: What is the equation linking ΔHsol, ΔHlatt, and ΔHhyd?

A. ΔHsol = ΔHlatt + ΔHhyd
B. ΔHsol = ΔHhyd – ΔHlatt
C. ΔHsol = –ΔHlatt + ΣΔHhyd
D. ΔHsol = ΔHlatt – ΣΔHhyd

Question 14: Which is the correct process for ΔHsol?

A. Salt + water → hydrated ions
B. Aqueous ions → gas
C. Solid → gas
D. Ion → electron + proton

Question 15: Which will have the most negative ΔHhyd?

A. Na⁺
B. Mg²⁺
C. K⁺
D. Ca²⁺

Question 16: What is entropy a measure of?

A. Energy stored in bonds
B. Randomness or disorder
C. Molecular mass
D. Concentration

Question 17: Which process leads to a decrease in entropy?

A. Melting
B. Evaporation
C. Freezing
D. Boiling

Question 18: What is ΔS if disorder increases?

A. Negative
B. Zero
C. Positive
D. Cannot determine

Question 19: What are the units of entropy (S)?

A. kJ mol⁻¹
B. kJ mol⁻¹ K⁻¹
C. J mol⁻¹ K⁻¹
D. J K mol⁻¹

Question 20: What is the sign of ΔS when a gas changes to a liquid?

A. Positive
B. Negative
C. Zero
D. Undefined

Question 21: What is the Gibbs free energy equation?

A. ΔG = ΔH × ΔS
B. ΔG = ΔH – TΔS
C. ΔG = TΔS – ΔH
D. ΔG = ΔH + TΔS

Question 22: For a reaction to be feasible at all temperatures, which combination is correct?

A. ΔH – and ΔS –
B. ΔH + and ΔS –
C. ΔH – and ΔS +
D. ΔH + and ΔS +

Question 23: Which combination is never feasible regardless of temperature?

A. ΔH – and ΔS –
B. ΔH + and ΔS –
C. ΔH + and ΔS +
D. ΔH – and ΔS +

Question 24: What does ΔG = 0 indicate?

A. Endothermic reaction
B. Reaction at equilibrium
C. Reaction is spontaneous
D. Reaction is irreversible

Question 25: Which of the following will increase entropy?

A. Liquid to solid
B. Decreasing number of gas molecules
C. Dissolving salt in water
D. Cooling a substance

Question 26: A reaction has ΔH = –100 kJ mol⁻¹ and ΔS = –200 J mol⁻¹ K⁻¹. At what temperature does it become non-feasible?

A. Always feasible
B. 500 K
C. 100 K
D. Never feasible

Question 27: Why is ΔS divided by 1000 in ΔG calculation?

A. To change units to Kelvin
B. To match ΔH units (kJ) with ΔS (J)
C. To convert ΔG to percentage
D. To eliminate T

Question 28: Which is most likely to have the highest standard entropy (S°)?

A. NaCl(s)
B. H₂O(l)
C. CO₂(g)
D. CH₄(g)

Question 29: What will be the sign of ΔS for 2SO₂(g) + O₂(g) → 2SO₃(g)?

A. Positive
B. Negative
C. Zero
D. Unpredictable

Question 30: What is the effect of increasing temperature on a reaction where ΔH is positive and ΔS is positive?

A. Less feasible
B. More feasible
C. No effect
D. Reversible

Answer Key and Detailed Explanations – A2 Level Chemistry: Chapter 23 Chemical Energetics Quiz

1. B. Energy released when gaseous ions form one mole of ionic solid
→ Lattice energy is the enthalpy change when one mole of solid ionic compound forms from its gaseous ions.

2. B. ½Cl₂(g) → Cl(g)
→ Atomisation is defined per mole of atoms; ½ mole of Cl₂ gives 1 mole of Cl atoms.

3. C. Lattice energy
→ The Born–Haber cycle is used to calculate lattice energy indirectly using Hess’s Law.

4. C. Smaller ionic radius and higher ionic charge
→ Greater electrostatic attraction gives more exothermic lattice energy.

5. C. Lattice formation
→ It is exothermic since stable ionic bonds are formed.

6. C. Cl(g) + e⁻ → Cl⁻(g)
→ Electron affinity is the energy change when an atom gains an electron.

7. B. Energy is needed to overcome repulsion from negative ion
→ Adding a second electron to already negative O⁻ is unfavorable due to repulsion.

8. B. Mg²⁺ and O²⁻
→ High charges and small sizes → stronger ionic attraction → more exothermic.

9. C. Atomisation of elements
→ Required step in the Born–Haber cycle to convert elements into gaseous atoms.

10. C. Ionisation energy
→ Removing electrons (especially the first few) is highly endothermic.

11. B. Energy released when gaseous ions become aqueous
→ Hydration enthalpy describes ion–water interactions.

12. C. Higher charge density
→ Small, highly charged ions are more strongly attracted to water molecules.

13. C. ΔHsol = –ΔHlatt + ΣΔHhyd
→ ΔHsol is determined by breaking lattice (endothermic) and hydrating ions (exothermic).

14. A. Salt + water → hydrated ions
→ Describes the process of solvation/dissolution.

15. B. Mg²⁺
→ Highest charge and smallest radius → most negative (exothermic) ΔHhyd.

16. B. Randomness or disorder
→ Entropy quantifies how dispersed energy or matter is in a system.

17. C. Freezing
→ Transition to a more ordered (lower entropy) state.

18. C. Positive
→ Disorder increases → positive entropy change.

19. C. J mol⁻¹ K⁻¹
→ SI unit for entropy.

20. B. Negative
→ Gas to liquid transition reduces disorder.

21. B. ΔG = ΔH – TΔS
→ Gibbs free energy formula (T in Kelvin, ΔS in J mol⁻¹ K⁻¹).

22. C. ΔH – and ΔS +
→ Always leads to negative ΔG, hence feasible at all T.

23. B. ΔH + and ΔS –
→ ΔG will always be positive → never feasible.

24. B. Reaction at equilibrium
→ ΔG = 0 implies system is at equilibrium.

25. C. Dissolving salt in water
→ Increases particle movement and disorder → positive ΔS.

26. B. 500 K
→ ΔG = 0 → T = ΔH / ΔS = 100,000 J / 200 J = 500 K.

27. B. To match ΔH units (kJ) with ΔS (J)
→ Ensures both energy terms in ΔG equation are in kJ.

28. C. CO₂(g)
→ Gases have higher entropy than liquids or solids.

29. B. Negative
→ 3 mol gas → 2 mol gas → reduction in disorder.

30. B. More feasible
→ ΔH + and ΔS + → increasing T favors negative ΔG.