group-II and Group IV of periodic table

Group–II A

1.      Alkaline earth metals do not exists free in nature.

2.      Magnesium and calcium are very abundant in the rocks of earth’s crust.

Magnesium Sources;

3.      Sea water

4.      Underground brines

5.      Mineral dolomite

6.      Magnesite (MgCO₃)

Calcium Sources;

7.      Sea shell (CaCO₃)

8.      Gypsum (CaSO₄.2H₂O)

Atomic and physical properties;

9.      All alkaline earth metal except Be are white in colour.

10. Alkaline earth metals are quite reactive and tarnish in air.

11. The value of their densities, melting point and boiling point are higher than those of alkali metals.

12. Down the group;

·        Atomic radius → Increases

·        Ionization energy → Decreases (exc;  Ra > Ba)

·        Electronegativity → Decreases (Ca=Mg and Ba=Ra)

·        Melting point and Boiling point → no regular term but greater than group first

·        Reactivity with water → Increases

13. Be does not react with water or steam at red heat.

Reaction with oxygen and Nitrogen;

14.  

	Simple oxide	Per oxide
Be	→	-
Mg	→	-
Ca	→	-
Sr	→	→
Ba	→	→
Ra	→	-

15. The nitrides of alkaline earth metals are ionic in nature except that of Ba which is covalent and unpredictable.

16. The reaction of alkaline earth metals with air rather than oxygen us complicated by the fact that  they all react with nitrogen to produce nitrides.

17. Mg + N₂ → Mg₃N₂

18. Be + N₂ → Be₃ N₂

19. If alkaline earth metals react with air, they form metal nitrides and metal oxides.

Trends in solubility of hydroxides, sulphates and carbonates

20. Solubility of hydroxides increases in water from top to bottom.

21. Solubility of sulphates decreases down the group.

22. Stability of carbonates increases down the group.

23. All carbonates of alkaline earth metals are insoluble in neutral medium while all dissolves in solids and decomposes at red heat.

24. CaSO₄ is sufficiently soluble in water.

25. Strontium and barium sulphates are almost insoluble.

26. Both carbonates and nitrates become more thermally stable as we go down the group of alkaline earth metals.

27. The one at lower position have to be heated more strongly than those at the top before they decompose.

Group–IV A 

28. C, Si, Ge, Sn and Pb.

29. M.P and B.P decreases down the group.

30. Melting point decreases because weaker bonds decreases with increase in atomic size.

31. Tin and Lid do not use all the four electrons for metallic bond.

32. C and Si are non-metals.

33. Ge is mettaloid.

34. Sn and Pb is metallic character.

35. Carbon and silicon show +4 oxidation state in carbonates and silicates.

36. Ge, Sn and Pb can show +2 and +4 oxidation states.

37. Oxidation state is defined as the apparent charge positive or negative on an atom of an element in a molecular ion.

38. Down the group, there is tendency for the sp² pair not to be used in bonding. This is called inert pair effect which dominates in lead.

39. Fajan’s rule; Sn⁺⁴ is smaller than Sn⁺² so the compounds of Sn⁺⁴ are covalent, while those of Sn⁺² are ionic.

40. All these elements give tetrachlorides (MCl₄) which are covalent and tetra-hedral due to sp³ hybrid orbitals.

41. The stability decreases from CCl₄ to PbCl₄.

42. PbCl₄ decreases to give PbCl₂ and Cl₂ gas.

43. At the top of the group, most stable oxidation state is +4 as shown as shown by carbon and silicon in CCl₄ and SiCL₄, they have no tendency to form dichlorides.

Reaction with water

44. CCl₄ does not react with water due to bulky nature of chlorine atoms around small carbon atom.

45. SiCl₄ and PbCl₄ react violently with water to reduce their respective oxides and fumes of HCl.

46. SiCL₄ + 2H₂O → SiO₂ +4HCl

47. SiO₂ → White

48. PbO₂ → Brown

49. PbCl₄ → covalent

50. PbCl₂ → ionic

51. PbCl₂ is sparingly soluble in cold water but more soluble in hot water.

Oxides

52. The elements of group-IV form two types of oxides i.e. monoxide and dioxide.

53. Monoxide include;

·        C O

·        Sn O

·        Pb O

54. Dioxide include;

·        C O₂

·        Sn O₂

·        Pb O₂

55. Non-metal oxidizes covalent in nature, such as oxidize of carbon and silicon.

56. Metal oxidize are ionic and nature such as oxidize of tin and lead.

57. CO₂ → Gas

58. Si O₂ → Solid

59. The SiO₂ is a giant covalent structure in which each silicon atom is bonded to four oxygen atoms through single covalent bonds whereas each oxygen atoms is bonded to two silicon atoms.

60. The geometry of the SiO₂ is tetrahedral (diamond like).

Acid Base Behavior of Group IV oxidizes

The acidity if group-IV oxides decreases as we go down the group.

61. C O₂    → acidic

62. Si O₂   → acidic

63. Ge O₂ → amphoteric

64. Sn O₂  → amphoteric

65. Pb O₂  → amphoteric

66. C O → neutral

67. Sn O → amphoteric

68. Pb O → amphoteric