3. Periodic Classification of Elements

 2 & 3 Marks

1. Write a note on Döbereiner Triads.

2. Give Newlands’ law of Octaves. What is its drawback?

3. Define Periodic law.

4. What are the Advantages of Mendeleev’s Periodic Table?

5. What are the Anomalies of Mendeleev’s Periodic Table?

6. Explain the work of Mosley and modern periodic law.

7. Define Modern periodic law.

8. What is the basic difference in approach between Mendeleev's periodic table and Modern periodic table?

9. What is periodicity?

10. How are properties of the elements correlated to the Electronic configuration?

11. What are Groups and Periods?

12. Elements a, b, c and d have the following Electronic configurations:

      a) 1s2, 2s2, 2p6      b) 1s2, 2s2, 2p6, 3s2, 3p1       c) 1s2, 2s2, 2p6, 3s2, 3p6      d) 1s2, 2s2, 2p1

      Which elements among these will belong to the same group of periodic table?

13. Name of elements with Atomic number above / greater than 100 using IUPAC Nomenclature.

14. In what Period and Group will an element with Z = 118 will be present?

15. The element with Atomic number 120 has not been discovered so far. What would be the IUPAC Name and the symbol for this element? Predict the possible Electronic configuration of this element.

16. Justify that the fifth period of the periodic table should have 18 elements on the basis of Quantum

      numbers.

17. What are s - block elements?

18. What are p - block elements?

19. What are d - block elements?

20. What are f - block elements?

21. Give the General electronic configuration of Lanthanides and Actinides?

22. Predict the position of the element in periodic table satisfying the Electronic configuration (n – 1)d2, ns2 where n = 5.

23. Define Atomic radius.

24. What is Covalent radius?

25. Covalent radius is always shorter than the actual atomic radius. Why?

26. Write Schomaker and Stevenson Equation and explain the terms.

27. Define Metallic radius.

28. Explain Variation of the Atomic Radius in Periods.

29. What is Screening effect?

30. What is Effective Nuclear Charge?

31. Using Slater's rule calculate the Effective nuclear charge on a 3p electron in Aluminium and Chlorine. Explain how these results relate to the atomic radii of the two atoms.

32. Explain Variation of the Atomic Radius in Groups.

33. Define Ionic radius.

34. Define Ionisation enthalpy (Ionisation energy). Give its unit.

35. Define Second ionisation energy.

36. Define Third ionisation energy.

37. Define Fourth ionisation energy.

38. The Effective nuclear charge of the cation is higher than the corresponding neutral atom. Why?

39. Explain Variation of the Ionisation energy in Periods.

40. Boron has higher Ionisation energy than Beryllium. Why?

41. Nitrogen has higher Ionisation energy than Oxygen. Why? Or Explain the following, give appropriate reasons – Ionisation potential of N is greater than that of O.

42. Explain Variation of the Ionisation energy in Groups.

43. Is the definition given below for Ionisation enthalpy is correct?

      "Ionisation enthalpy is defined as the energy required to remove the most loosely bound electron from the valence shell of an atom".

44. Magnesium loses electrons successively to form Mg+, Mg2+ and Mg3+ ions. Which step will have the highest Ionisation energy and why?

45. How would you explain the fact that the Second Ionisation potential is always higher than First Ionisation potential?

46. The electronic configuration of atom is one of the important factor which affects the value of Ionisation potential and Electron gain enthalpy. Explain.

47. Why the First Ionisation enthalpy of Sodium is lower than that of Magnesium while it’s Second  

      Ionisation enthalpy is higher than that of Magnesium?

48. Explain the following, give appropriate reasons – First Ionisation potential of C -atom is greater than that of B atom, where as the reverse is true is for Second Ionisation potential.

49. Define Electron gain enthalpy (Electron affinity).

50. Explain the following, give appropriate reasons – The Electron affinity values of Be, Mg and Noble gases are Zero and those of N (0.02 eV) and P (0.80 eV) are Very Low.

51. Explain the following, give appropriate reasons – The formation of F–(g) from F(g) is Exothermic while that of O2–(g) from O(g) is Endothermic.

52. Explain Variation of the Electron affinity in Periods.

53. Explain Variation of the Electron affinity in Groups.

54. Define Electronegativity.

55. Briefly give the basis for Pauling's scale of electronegativity.

56. What is Valence of an atom?

57. Elements of extreme ends of the periodic table show high reactivity when compared to the elements present in the middle. Why?

58. The Noble gases are chemically inert in nature. Why?

59. Explain how the Ionisation energy related to the metallic and non-metallic character.

60. Why halogens act as oxidising agents?

61. What are Isoelectronic ions? Give examples.

62. Mention / Explain the Anomalous properties of second period elements.

63. Explain the Diagonal relationship.

64. The hydroxides of Alkali metals become more basic. Why?

65. Prove – Beryllium hydroxide is amphoteric in nature.

      Problem

      Covalent radius, Effective nuclear (Slater's rule), Ionic radii (Pauling’s method), Ionisation enthalpy

5 Marks

1. Analyse the change in the electronic configuration of elements along the periods.

2. Analyse the change in the electronic configuration of elements down the groups.

3. Explain the periodic trend of Atomic Radius

4. Describe about Slater's rule.

5. Explain the Pauling method for the determination of Ionic radius.

6. Explain the periodic trend of Ionisation potential.

7. Explain the periodic trend of Electron affinity / Electron gain enthalpy.

8. State the trends in the variation of Electronegativity in groups and periods.