Level 1 — Basic Recall (1–10)

1. Define enthalpy.

2. What is an exothermic reaction?

3. What is an endothermic reaction?

4. State the sign of ΔH for exothermic reactions.

5. State the sign of ΔH for endothermic reactions.

6. Define activation energy.

7. What is the standard enthalpy of combustion?

8. Define the standard enthalpy of formation.

9. What is bond energy?

10. Draw a simple energy profile diagram for an exothermic reaction.

Level 2 — Understanding (11–20)

11. Explain why bond breaking is endothermic.

12. Explain why bond forming is exothermic.

13. Why do exothermic reactions feel hot?

14. Why do endothermic reactions feel cold?

15. Describe how a catalyst affects activation energy.

16. Explain the significance of standard conditions.

17. Compare ΔHf° and ΔHc°.

18. Explain why ΔHneutralization is always exothermic for strong acids/bases.

19. Illustrate how enthalpy changes can be measured using calorimetry.

20. Explain the difference between ΔH and ΔU (enthalpy vs internal energy).

Level 3 — Application (21–30)

21. Calculate ΔH using values of bonds broken and formed.

22. Use Hess’s Law to determine ΔH when direct measurement is difficult.

23. Draw an enthalpy cycle for formation → combustion relationship.

24. Calculate heat change using q = mcΔT.

25. Explain why combustion reactions are always exothermic.

26. Calculate ΔHreaction using enthalpies of formation.

27. Predict whether a bond with high bond energy is strong or weak.

28. Use an energy profile to compare uncatalysed vs catalysed pathways.

29. Explain why dissolving ammonium nitrate in water is endothermic.

30. Determine ΔH for a reaction using Hess’s cycle with two intermediate steps.

Level 4 — Analysis (31–40)

31. Analyse why breaking strong bonds requires more energy.

32. Compare the enthalpy change of neutralization for strong vs weak acids.

33. Evaluate errors in calorimetry experiments.

34. Explain why enthalpy alone cannot determine spontaneity.

35. Discuss why some reactions with positive ΔH still occur spontaneously.

36. Explain how entropy influences reaction feasibility.

37. Analyse the relationship ΔG = ΔH – TΔS.

38. Predict spontaneity at high vs low temperatures.

39. Compare entropy in solids, liquids, and gases.

40. Explain why ΔS increases when a gas is produced.

Level 5 — Exam/Challenge (41–50)

41. Use Gibbs free energy to predict spontaneity under specific conditions.

42. Determine the temperature at which a reaction becomes spontaneous (ΔG = 0).

43. Analyse how entropy affects phase changes (melting, boiling).

44. Calculate ΔG using values for ΔH and ΔS.

45. Discuss the thermodynamic reasoning behind dissolving salts (entropy-driven).

46. Derive Hess’s Law using the First Law of Thermodynamics.

47. Evaluate the feasibility of industrial reactions based on ΔG.

48. Compare lattice enthalpies of ionic crystals using size and charge.

49. Analyse the thermodynamic stability of compounds using enthalpy cycles.

Explain why exothermic reactions are often fast but endothermic ones slow—and note exceptions using activation energy.