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