Level 1 — Basic Recall (1–10)

1. What is meant by electronic configuration?

2. State the maximum number of electrons in the 1st shell.

3. State the maximum number of electrons in any orbital.

4. What letter symbols represent the four subshells?

5. Define an orbital.

6. What does Pauli’s exclusion principle state?

7. State Hund’s rule.

8. What is the Aufbau principle?

9. Which subshell fills first: 4s or 3d?

10. What is the full electron configuration of oxygen?

Level 2 — Understanding (11–20)

11. Why do 4s orbitals fill before 3d?

12. Explain why electrons prefer to occupy separate orbitals before pairing.

13. What is meant by “valence electrons”?

14. Define atomic radius.

15. Define ionization energy.

16. What is electronegativity?

17. Describe how atomic radius changes across a period.

18. Explain the trend in ionization energy across a period.

19. Why does atomic radius increase down a group?

20. Why is fluorine the most electronegative element?

Level 3 — Application (21–30)

21. Write the electron configuration of Ca²⁺.

22. Write the shorthand configuration of chlorine.

23. Predict the period and group of an element with configuration [Ne]3s²3p⁵.

24. Explain why sodium forms a +1 ion.

25. Compare the atomic radius of Na and Mg.

26. Arrange F, Cl, Br in increasing atomic size.

27. Explain why ionization energy of sulfur is slightly lower than that of phosphorus.

28. Predict which element is more metallic: Mg or Sr.

29. Show electron configuration for Cu and explain its irregularity.

30. Predict the number of unpaired electrons in nitrogen.

Level 4 — Analysis (31–40)

31. Analyse why first ionization energy increases across a period.

32. Explain the drop in ionization energy from Be to B.

33. Explain the drop in ionization energy from N to O.

34. Compare the shielding effect in Li vs Cs.

35. Analyse why transition metals form colored ions in terms of electron configuration.

36. Evaluate why noble gases have the highest ionization energies.

37. Compare Na and Na⁺ in terms of radius.

38. Predict trends in metallic character and justify.

39. Explain how d-block contraction affects atomic properties.

40. Analyse why electron affinity becomes less exothermic down a group.

Level 5 — Exam/Challenge (41–50)

41. Explain the periodic trends using effective nuclear charge.

42. Predict and explain irregularities in Period 3 ionization energies.

43. Compare electron configurations of Fe²⁺ and Fe³⁺.

44. Explain the reasoning behind f-block filling order.

45. Use quantum numbers to describe a specific electron in phosphorus.

46. Analyse the periodic pattern in successive ionization energies of magnesium.

47. Discuss the role of electron-electron repulsion in orbital energy levels.

48. Predict chemical behavior based on electron configuration.

49. Explain why 4s electrons are removed first during ionization of transition metals.

50. Evaluate how periodicity supports the modern arrangement of the periodic table.