1. p-block elements
Grade 12

2. Nitrogen Family Electronic configuration- ns2np3
Atomic and ionic radii-
Increases as we move down the group. But there is small increase in the size from As to Bi due to completely filled d- and f- orbitals.
Q. Atomic radius of elements increases in a group, but there is only a small increase from As to Bi.
A. This is due to the presence of completely filled d-orbitals and/or f-orbitals in them, which have weak screening effect and more effective nuclear charge called as inert- pair effect.
Ionization Enthalpy-
Decreases down the group.
Q. I.E. of Group 15 elements is much greater than that of the Group 14 and Group 16 elements in the corresponding periods.
A. Because of the extra stability of a half-filled p-orbital configuration.
Electronegativity-

3. Chemical Properties
Oxidation States and Trends in Chemical Reactivity-
Common oxidation states are -3,+3 and +5.
Q. Nitrogen does not exhibit covalency greater than 4.
A. Due to the absence of d-orbitals in the valence shell of N.
2. Tendency to exhibit -3 oxidation state decreases down the group due to increase in size and metallic character.

4. The stability of +5 oxidation state decreases down the group but stability of +3 oxidation state increases due to inert pair effect.
Q. The stability of +5 state decreases and that of +3 state increases down the group.
A. Due to inert pair effect. (explain inert pair effect)
Q. Metallic character increases down the group.
A. Due to decrease in I.E. and increase in atomic size.

5. Anomalous properties of Nitrogen
Nitrogen has small size, high electronegativity,high IP and non-availability of d-orbitals.
Form p-p pie bonds with itself.
Nitrogen exists as diatomic with one triple bond.
Q. Nitrogen exists as diatomic molecule N2 gas at room tempertaure.
A. Due to the ability of the N atom to form pπ-pπ triple bond with another N atom.
Bond enthalpy is very high.
P-P, As-As, Sb-Sb have single bonds whereas Bi forms metallic bond.

6. Q. Nitrogen shows anomalous properties.
A. Due to its small size, high electronegativity, high I.E., non-availability of d-orbitals in its valence shell and its ability to form pπ-pπ multiple bonds with itself, and with other elements like C and O.

7. Weak catenation tendency.
Q. Tendency for catenation is less in N as compared to the other members of the group.
A. N-N single bond is weaker than those of the other elements. This is due to the high inter-electronic repulsion of the non-bonding electrons, owing to the small N-N bond length.
Non-availability of d- orbitals.
Reactivity towards hydrogen-All elements form hydrides of EH3 type. Stability of Hydrides decreases from NH3 to BiH3, while the reducing character increases. Basicity trend: NH3>PH3>AsH3>SbH3>BiH3
Q. Thermal stability of hydrides decreases from NH3 to BiH3.
E-H bond enthalpy (bond strength) decreases down the group. (E  Group 15 element)
Q. Reducing character of EH3 increases from NH3 to BiH3.
A. E-H bond strength decreases, so H is easily supplied for reduction, on moving down the group.

8. Q. Basic character decreases down the group for EH3.
A. Due to the presence of a lone pair of electrons on the central atom, they behave as Lewis bases. N being the smallest in size, electron-density on N is high. Thus, it is the strongest base. Base strength decreases with increase in the size of the central atom.
Q. NH3 is a stronger base than PH3. (phosphine).
A. Due to smaller size of N, electron-density on N is higher. Thus, it is a strong Lewis base.
Q. NH3 has exceptionally high b.p.
A. Due to the high electronegativity of N, intermolecular H bonds exist between NH3 molecules.

9. Q. Among hydrides of Group 15 elements, BiH3 is the strongest reducing agent.
A. BiH3 is the least stable, because Bi-H bond is the weakest (least bond enthalpy) compared to the E-H bond of the other elements of the group.
Q. NH3 molecule is triagonal pyramidal in shape.
A. Due to the presence of 3 bp and 1 lp of electrons on the central N atom in the molecule.Structure.
Q. NH3 is used in the detection of metal ions such as Cu2+, Ag+ etc.
A. Due to the presence of a lone pair on the central N atom, NH3 is a Lewis base. Thus, it donates electron pairs to metal ions and hence forms coloured complexes which help in the detection of metal ions.
e.g Cu2+(aq) NH3 → [Cu(NH3)4]2+(aq)
deep blue solution

10. 2. Reactivity towards oxygen-
Two types of oxides : E2O3 and E2O5 type.
The oxide in higher oxidation state are more acidic than the lower ones. Acidic character decreases down the group.N2O3 and P2O3 are acidic,As2O3 and Sb2O3 are amphoteric and Bi2O3 are basic.

11. 3. Reactivity towards halogens-
Two types-EX3 and EX5.Nitrogen do not form pentahalides
due to the non-availability of d-orbitals. Pentahalides are
more covalent than trihalides.
Q. PCl5 exists, but NCl5 does not.
A. Due to the absence of d-orbitals in the valence shell of N,N cannot expand its valence shell.
Q. Pentahalides of Group 15 elements are more covalent than trihalides.
A. Higher the positive oxidation state of the central atom, the more is its polarizing power and, therefore, more the covalent nature of the bond.
4. Reactivity towards metals-
Form binary compounds exhibiting -3 oxidation states.e.g. Calcium nitride, Calcium phosphide etc.

12. Dinitrogen Preparation-
In the laboratory, nitrogen is prepared by heating a mixture of ammonium chloride and sodium nitrite and a small quantity of water. If ammonium nitrite is heated by itself it decomposes to produce nitrogen gas. However, this reaction is very fast and may prove to be explosive.
                                               

13. By thermal decomposition of ammonium dichromate-
(NH4)2Cr2O N2+4H2O+Cr2O3
3.By thermal decomposition of sodium or barium azide-
Ba(N3) Ba+3N2

14. Properties of Dinitrogen
Dinitrogen is inert at room temperature due to the high bond enthalpy of N-N triple bond.
Q. N2 gas is chemically inert at room temperature.
A. Due to very high bond enthalpy of N≡N.
2. At high temperature reactivity increases and nitrogen form ionic nitrides with metals and covalent nitrides with non-metals.
6Li+N Li3N

15. 3Mg+N Mg3N2
It reacts with hydrogen at 773 K in the presence of catalyst forms ammonia (Haber`s Process).
N2+3H NH3
It reacts with dioxygen at very high temperature and forms nitric oxide NO.
N2+O NO

16. Ammonia Preparation- By decay of nitrogenous organic matter-
NH2CONH2+2H2O (NH4)2CO3
2NH3+CO2

17. On small scale ammonia can be obtained by decomposition of ammonium salts with caustic soda or lime.
NH4Cl+Ca(OH) NH3+H2O+CaCl2
On large scale ammonia can be manufactured by Haber`s process.
N2+3H NH3

18. Properties of ammonia
Has high melting and boiling point due to H2 bonding.
Trigonal pyramidal in shape.
Ammonia is highly soluble in water.Its aqueous solution gives OH- in aqueous solution.
NH3+H2O NH4+OH-
4. As a weak base, it precipitates the hydroxides of many metals from their salt solutions.

19. 2FeCl3+3NH4OH Fe2O3.H2O+3NH4Cl
(brown ppt.)
5. Acts as Lewis base due to the presence of lone pair of electron on N2.e.g.
Cu+2+4NH [Cu(NH3)4] +2
(deep blue)

20. Oxides of Nitrogen N2O NH4NO3 (s) → 2 H2O (g) + N2O (g)
Nitrous oxide
                      
                          
                                     
Oxides of Nitrogen N2O
IUPAC name Dinitrogen monoxide Other name Laughing gas
Structure-
NH4NO3 (s) → 2 H2O (g) + N2O (g)

21. NO
Preparation NaNO2+2 FeSO4+3 H2SO4 → Fe2(SO4)3+2 NaHSO4 + 2 H2O + 2 NO Preferred IUPAC name Nitric oxide Systematic name Nitrogen monoxide
Other names Nitrogen(II) oxide
               

22. N2O3 Dinitrogen trioxide
                              
N2O3
Dinitrogen trioxide
It forms upon mixing equal parts of nitric oxide and nitrogen dioxide and cooling the mixture below −21 °C (−6 °F):
NO + NO2    N2O3

23. The thermal decomposition of some metal nitrates also affords NO2:
Nitrogen dioxide
Nitrogen dioxide
                      
The thermal decomposition of some metal nitrates also affords NO2:
2 Pb(NO3)2 → 2 PbO + 4 NO2 + O2

24. Dinitrogen tetroxide
                              
N2O4
Dinitrogen tetroxide
N2O4 ⇌ 2 NO2

25. Q. NO2 dimerises.
A. NO2 contains an odd no. of valence electrons. Thus, it behaves as a typical odd molecule. On dimerisation, it is converted into a more stable N2O4 molecule with an even no. of electrons.

26. N2O5 P4O10 + 12 HNO3 → 4 H3PO4 + 6 N2O5 Dinitrogen pentoxide
                              
Dinitrogen pentoxide
                              
N2O5
Dinitrogen pentoxide
P4O HNO3 → 4 H3PO4 + 6 N2O5

27. Nitric Acid
1. Nitric acid is made by reacting nitrogen dioxide (NO2) with water.
3 NO2 + H2O → 2 HNO3 +NO
2. Almost pure nitric acid can be made by adding sulfuric acid to a nitrate salt, and heating the mixture with an oil bath. A condenser is used to condense the nitric acid fumes that bubble out of the solution.
2NaNO3 + H2SO HNO3+Na2SO4

28. On a large scale it can be manufactured by Ostwald`s method
On a large scale it can be manufactured by Ostwald`s method.By catalytic (Pt/Rh gauge) oxidation of ammonia with atmospheric oxygen.
4NH3+5O NO+6H2O
2NO+O NO2
3NO2+H2O HNO3+NO

29. Properties of HNO3
Nitric acid
               
                                 
                      
Nitric acid

30. Chemical Properties of HNO3
Nitric acid is a strong oxidizing agent. When it undergoes thermal decomposition it yields nascent oxygen as follows:
2HNO H2O+2NO2+[O]
Remember:-
The nascent oxygen so formed oxidizes non-metals, metals, inorganic as well as organic compounds etc.

31. a) With Non-Metals
With hot concentrated nitric acid, non-metals are oxidized to their oxide while the acid itself gets reduced to nitrogen dioxide.
In all the reactions described below, the nascent oxygen is released during the thermal decomposition of the acid, which oxidizes the non-metals.
 2HNO H2O+2NO2+[O]
i) With carbon
Nascent oxygen reacts with carbon to form carbon dioxide.
C+2[O] CO2
The overall reaction is
4HNO3+C H2O+4NO2+CO2

32. ii) With sulphur
Nascent oxygen reacts with sulphur to form sulphur trioxide.
S+3[O] SO3
Sulphur trioxide in turn reacts with water to form sulphuric acid
H2O+SO H2SO4
The overall reactions is:
S+6HNO H2O+H2SO4+6NO2

33. iii) With phosphorous
Nascent oxygen reacts with phosphorous to form phosphorous (V) oxide.
P4+10[O] P4O10
Phosphorous (V) oxide in turn reacts with water to form phosphoric acid.
6H20+P4O H3PO4
The overall reaction is:
P4+20HNO H2O+4H3PO4+20NO2

34. b) With Metals
Nitric acid behaves differently with different metals at different concentrations.
i) With sodium, potassium and calcium the reaction is highly explosive.
ii) With magnesium and manganese
With magnesium and manganese, cold and extremely dilute (1%) nitric acid, reacts to yield hydrogen.

35. iii) With Copper
With cold dilute nitric acid:
Copper reacts with cold and dilute nitric acid to yield copper nitrate, water and nitric oxide.
3Cu+8HNO Cu(NO3)2+4H2O+2NO
The formed nitric oxide combines with the oxygen of air to give brown fumes of NO2.
2NO+O NO2
With concentrated nitric acid (cold or hot):
Copper reacts with cold or hot concentrated nitric acid to yield copper nitrate, water and nitrogen dioxide.
Cu+4HNO Cu(NO3)2+2H2O+2NO2

36. iv) With Zinc
With cold dilute nitric acid:
Zinc reacts with cold and dilute nitric acid to yield zinc nitrate, water and nitric oxide.
  3Zn+8HNO Zn(NO3)2+4H2O+2N2O
With concentrated nitric acid (cold or hot):
Zinc reacts with cold or hot concentrated nitric acid to yield zinc nitrate, water and nitrogen dioxide.
Zn+4HNO Zn(NO3)2+2H2O+2NO2

37. Some metals like Cr, Al do not oxidise in conc
Some metals like Cr, Al do not oxidise in conc. Nitric acid because of the formation of a thin film of oxide on its surface.
Conc. Nitric acid also oxidises non-metals.
I2+10 HNO HIO3+10NO2+4H2O
C+4HNO CO2+2H2O+2NO2
S8+48HNO H2SO4+48NO2+16H2O
P4+20HNO H3PO4+20NO2+4H2O

38. Q. Conc. nitric acid is a strong oxidizing agent and oxidises metals like Cu and Zn. However, metals like Cr and Al do not dissolve in conc. HNO3.
A. Due to the formation of passive oxide films on the surface of Cr and Al.

39. Phosphorus-Allotropic Forms White Phosphorus
Translucent, white, waxy solid.
Insoluble in water but soluble in CS2.
Dissolves in boiling NaOH solution.
P4+3NaOH+3H2O PH3+3NaH2PO2
It ignites spontaneously in air at about 50C and at much lower temp. if finely divided. This combustion gives phosphorus (V) oxide:P4+5O2 P4O10

40. Red Phosphorus
It is prepared by heating white phosphorus to about 540 K in an inert atmosphere of nitrogen for several hours which on heating and under high pressure gives black phosphorus.
(i) It is a hard crystalline solid without any smell and is poisonous in nature.
(ii) It is insoluble in water as well as in carbon disulphide.
(iii) It is more stable and relatively less reactive.
(iv) It consists of tetrahedral units of P4 linked to one another to constitute linear chains.

41. Black Phosphorus
It is prepared by heating white phosphorus to about 470 K under high pressure of 1200 atmospheres in inert atmosphere. 
(i) It has metallic lustre.
(ii) It is most inactive form of phosphorus.
(iii) It has a layer type structure in which each layer consists of phosphorus atoms.

42. Phosphine Preparation- Reaction of Calcium phosphide with
water or dil. acid.
Ca3P2+6H2O 3Ca(OH)2+2PH3
Ca3P2+6HCl 3CaCl2+2PH3
2. Lab. Method- Heating white phosphorus with conc. NaOH in inert atmosphere of CO2.

43. P4+3NaOH+3H2O PH3+3NaH2PO2
3. From impure phosphine-
PH3+HI PH4I (+KOH) KI+H2O+PH3
Properties-
Slightly soluble in water.
Aqueous solution of phosphine decomposes in light and gives red phosphorus and hydrogen.

44. 3CuSO4+2PH Cu3P2+3H2SO4
3HgCl2+2PH Hg3P2+3HCl
PH3+HBr PH4Br
Q. PH3 is a Lewis base.
A. Due to the presence of a lone pair of electrons on the P atom.
Q. Bond angle in PH4+ is higher than that of PH3.
A. In PH4+, there are 4 bond pairs on the central P atom. Thus, it has a tetrahedral shape with a bond angle of 109.5. However, in PH3, there is a lone pair on the central P atom. Due to lp-bp repulsion, bond angle in PH3 is slightly less than 109.5.

45. Phosphorus Trichloride
Preparation-
Lab method-Heat white phosphorus in current of dry chlorine.
P4+6Cl PCl3
2. P4+8SOCl PCl3+4SO2+2S2Cl2

46. Properties-
PCl3+3H2O H3PO3+3HCl
PCl3+3CH3COOH H3PO3+CH3COCl
PCl3+3C2H5OH H3PO3+C2H5Cl
Structure-
sp3 hybridisation. Pyramidal structure.

47. Phosphorus Pentachloride
Preparation-
Reaction of white phosphorus with excess of dry chlorine.
P4+10Cl PCl5
P4+10SO2Cl PCl5+10SO2
Properties-
PCl5+H2O POCl3+2HCl
POCl3+3H2O H3PO4+3HCl

48. 2. On heating
PCl PCl3+Cl2
3. C2H5OH+PCl C2H5Cl+POCl3+HCl
4. 2Ag+PCl AgCl+PCl3
5. Sn+2PCl SnCl4+2PCl3
Structure-
In gaseous and liquid phases, it has a trigonal bipyramidal structure. 3 equatorial P-Cl bonds are equivalent while 2 axial bonds are longer than equatorial bonds.

49. Q. PCl3 fumes in the presence of moisture.
A. This is due to the formation of HCl.
PCl3 + 3H2O  H3PO4 + 3HCl
Q. All the 5 P-Cl bonds in PCl5 are not equivalent.
A. PCl5 molecule is trigonal bipyramidal in shape. The 3 equatorial bonds are of the same length but the 2 axial bonds are longer due to repulsion between bond pairs of the axial and equatorial bonds.
Q. PCl5 in the gaseous and solid states, respectively, do not have the same geometry.
A. PCl5 (g) is trigonal bipyramidal in shape. In the solid state, PCl5 is ionic with tetrahedral [PCl4]+ cation and octahedral [PCl6]- anion.

50. Oxoacids of Phosphorus
Oxidation State/ Formula /Name/ Acidic Protons/Compounds
+1/H3PO2 /hypophosphorous acid /1/ acid, salts
+3/H3PO3/(ortho)phosphorous acid/ 2/ acid, salts
+5/(HPO3)n/metaphosphoric acids/ n /salts (n=3,4)
+5/H5P3O10/triphosphoric acid /3/ salts
+5/H4P2O7/pyrophosphoric acid /4/ acid, salts
+5/H3PO4/(ortho)phosphoric acid/ 3/ acid, salts

52. Q. Orthophosphorous acid H3PO3 undergoes disproportionation.
A. In H3PO3, P is in the +3 state. ─3 and +5 states of P are more stable. So, it undergoes disproportionation.
4H3PO3  3H3PO4 + PH3
Q. Ortho phosphorous acid H3PO3 (phosphonic acid) is diprotic (or dibasic) while ortho phosphoric acid H3PO4 is triprotic (or tribasic).
A. Only those H atoms which are attached to oxygen in P-OH form are ionisable and cause basicity. Thus, H3PO3 with only 2 P-OH bonds is dibasic and H3PO4 with 3 P-OH bonds is tribasic.
Q. Hypophosphoric acid H3PO2 (phosphinic acid) is monobasic or monoprotic.
A. There is only one P-OH bond. Thus, there is only 1 ionisable H atom and, hence only one H+ ion can be liberated.
Q. H3PO2 is a strong reducing agent.
A. Oxoacids with P-H bonds have reducing properties and H3PO2 (above fig.) has 2 P-H bonds in it.

53. Q. NH3 forms hydrogen bonds but PH3 does not.
A. Due to the small size and high electronegativity of the central N atom, NH3 is able to form intermolecular H-bonds.
Q. R3P=O exists, but R3N=O does not. (R  alkyl group)
A. P can form dπ-pπ bond with O. However, N, due to the absence of d-orbitals, cannot form dπ-pπ bond with O.
Q. P(C2H5)3 and As(C6H5)3 act as ligands and bond with transition metals.
A. Due to the ability of P and As to form dπ-dπ bond with transition metals.
Q. Nitrogen exists as N2 gas while phosphorous exists as P4 solid.
A. N can form pπ-pπ triple bond with another N atom while P, due to its larger size, cannot form pπ-pπ bond with another P atom.
Q. The HNH bond angle in NH3 is higher than HPH, HAsH and HSbH bond angles in their respective hydrides.
A. Due to the small size of the central N atom and the small length of the N-H bond, bp-bp repulsion in NH3 is greater than that of the other halides of the group. Moreover, in NH3, hybridisation is sp3 while in heavier elements the bond is between s and p orbitals of hydrogen and the element respectively.

54. Oxygen Family Electronic Configuration-ns2np4.
Atomic and Ionic Radii- Increases as we move down the group.
Ionisation Enthalpy- Decreases down the group.
Melting and Boiling Point- Increases down the group.

55. Electron Gain Enthalpy- Value becomes less negative as we move down the group.
Q. Electron gain enthalpy of O is less negative than that of S.
A. Due to the very small size of the O atom there is greater inter-electronic repulsion in the relatively smaller 2p sub-shell hence the incoming electron experiences less attraction in oxygen atom.
Electronegativity- Decreases down the group.
Q. O2 exists as a gas while other members are solids.
A. Due to the ability of O to form pπ-pπ double bond with another O atom.

56. Oxidation States and Trends in Chemical Reactivity
Shows -2, -1 and +2 oxidation states.
Tendency to show -2 oxidation state decreases down the group.
Unlike oxygen, other elements show +2,+4 and +6 oxidation states.
Stability of +6 oxidation state decreases down the group but +4 increases due to inert – pair effect.
Q. O shows only negative oxidation state, except in OF2.
A. Since O is a highly electronegative element, it shows only ─2 oxidation state. However, in OF2, F is more electronegative due to which O has +2 oxidation state.

57. Q. Oxygen does not exhibit higher oxidation states (+4, +6) like the heavier members of its group.
A. Due to the absence of d-orbitals in the valence shell of oxygen.
Q. Stability of +6 state decreases and that of +4 state increases down Group 16.
A. Due to inert pair effect.

58. Anomalous Behaviour of Oxygen
Small size , high electronegativity.
Absence of d- orbital.
Presence of strong hydrogen bonding in H2O but absent in H2S.

59. Reactivity with Hydrogen RTA
Form hydrides of type H2E.
Acidic character increases from H2O to H2Te.A
Thermal stability decreases down the group.T
All hydrides except H2O possess reducing property and it increases from H2S to H2Te.R

60. Q. Acidic character of the hydrides of Group 16 increases down the group.
A. Due to decrease in H-E bond enthalpy. (E  Group 16 element)
Q. Thermal stability of H2E decreases down the group.
A. H-E bond strength decreases.
Q. Reducing character increases down the group.
A. H-E bond strength decreases.
Q. Water has exceptionally high m.p. and b.p.
A. Due to intermolecular H-bonds, owing to the high electronegativity and small size of the O atom.

61. Reactivity with Oxygen
Form EO2 and EO3 types of oxides.
Reducing property of dioxide decreases from SO2 to TeO2.
Both EO2 and EO3 types are acidic in nature.
Q. Reducing property of dioxides reduces from SO2 to TeO2.
A. Due to inert pair effect the stability of +6 oxidation state decreases from S to Te.

62. Reactivity towards Halogens
Form EX6, EX4 and EX2 types.
Stability of halides decreases from Fluoride to iodide.
Among EX6 type, hexafluorides are only stable.
SF6 is exceptionally stable for steric hindrance.

63. Q. Stability of halides of Group 16 decreases in the order F > Cl > Br > I.
A. Stability of halides decrease with a decrease in electronegativity of the halogen.
Q. SF6 is very stable, chemically inert and resistant to hydrolysis.
A. The S atom in SF6 is sterically protected by 6 F atoms.
Q. SF4 and SeF6 are more easily hydrolysed.
A. They are less sterically protected. In SF4 steric protection is less due to lesser no. of F atoms and in SeF6 steric protection is less due to larger size of Se.

64. Dioxygen Preparation- Lab method- 2KClO3 2KCl + 3O2(heat,MnO2)
By thermal decomposition-
2Ag2O Ag + O2
2HgO Hg+O2
2Pb3O PbO +O2

65. 2H2O H2O + O2 (MnO2)
Industrial method- From water or air. By electrolysis of water.
Properties-
1. 2Ca + O CaO
4Al + 3O Al2O3
P4 + 5O P4O10
C + O CO2

66. 2 ZnS + 3 O ZnO + 2 SO2
CH4 + 2 O CO2 + 2 H2O
2 SO2 + O SO3(V2O5)
4 HCl + O Cl2 + 2 H2O

67. Simple Oxides
A binary compound of oxygen with another element is called oxide.
Oxides may be simple or mixed.
Simple oxides may be acidic, basic or amphoteric.
Oxide + H2O Acid (Acidic Oxide)
Example- SO2, CO2

68. SO2 + H2O H2SO3
Non- metal oxides are acidic but oxides of metals in high oxidation states also have acidic character. (e.g. Mn2O7, CrO3)
Oxide + H2O Base (Basic Oxide)
Examples- Na2O, CaO, BaO.
CaO + H2O Ca(OH)2
Metallic oxides are basic.

69. Some metallic oxides shows both acidic as well as basic character called as Amphoteric oxides.
Some oxides are neither acidic nor basic called as Neutral oxides.

70. Ozone
Preparation-
When a slow dry stream of oxygen is passed through a silent electrical discharge conversion of oxygen to ozone occurs. The product is OZONISED OXYGEN.
3O O3 – Heat
This is Endothermic Reaction.

71. Properties-
Powerful oxidizing agent.
O O2 + O
Oxidizes PbS to PbSO4 and I- to I2.
PbS + 4 O PbSO4 +4 O2
2 I- +H2O +O OH- + I2 + O2
3. NO + O NO2 + O2

72. Q. Though the formation of oxides with elements is exothermic, some external heating is required to initiate the reaction.
A. Bond dissociation enthalpy of O=O is very high.
Q. The 2 O-O bond lengths in the ozone molecule are identical.
Ozone is a resonance hybrid of 2 canonical structures.
Q. O3 is a powerful oxidizing agent.
A. O3  O2 + O
The above reaction is highly exothermic (H negative and S positive, thus, G negative). The nascent O liberated makes O3 a powerful oxidizing agent.

73. Sulphur-Allotropic Forms
Main allotropic forms are-
Yellow rhombic sulphur
2. Monoclinic sulphur
Sulphur forms yellow, rhombic crystals out of 8-membered rings of sulphur atoms (S8).
Sulphur forms yellow, monoclinic, needle-like crystals out of 8-membered rings of sulphur atoms (S8).
Plastic sulphur is yellow and made up of long chains of sulphur atoms. It reverts to S8 rings in time.

74. Sulphur Dioxide Preparation- Sulphur is burnt in air. S + O2 SO2
2. Lab method- Treating SO3- with H2SO4.
SO H SO2 + H2O
Industrial- As a by- product of roasting of sulphide ores.
4 FeS O Fe2O3 + 8SO2

75. Properties-
Colourless gas with pungent smell.
Soluble in water.
SO2 + H2O H2SO3(sulphurous acid)
2 NaOH + SO Na2SO3 + H2O
Na2SO3 + H2O + SO NaHSO3
SO2 + Cl SO2Cl2(sulphuryl chloride)
( +nce of charcoal )
2SO2 + O SO3( V2O5,Oxidation)

76. Moist SO2 behaves as reducing agent.
Fe (iii) Fe (ii)
Decolorizes acidified KMnO4 solution.
2Fe3+ + SO2 + 2H2O Fe2+ + SO H+
5SO2 + 2 MnO H2O SO H+ + 2 Mn2+

77. Oxoacids of Sulphur

78. Sulphuric Acid Manufacture- By Contact Process
i) S8+8 O SO2(burning sulphur)
4FeS2+11 O Fe2O3 + 8 SO2 (roasting of sulphur rich ore)
ii) 2SO2 + O SO3+ HEAT(oxidn.) (V2O5 or Platinised asbestos)
iii) SO3 + H2SO4(conc.) H2S2O7
H2S2O7 +H2O H2SO4

79. Properties-
Dissociation-
H2SO H2O + SO3
Acidic Nature-
H2SO H+ + HSO4-
HSO H+ + SO42-

80. As dehydrating agent-
i) Drying of gases-Gases which do not react with H2SO4 like CO2,SO2,Cl2.
ii) Charring-
C12H22O H2O + 12 C (conc.H2SO4)
Oxidizing Action-Strong oxidising agent.
Cu + 2H2SO CuSO4 + SO2 +2H20(conc.)

81. Halogen Family Electronic Configuration- ns2np5
Atomic and Ionic Radii- Increases as we move down the group.
Ionisation Enthalpy- Decreases down the group.
Electron Gain Enthalpy- Halogens have max. –ve electron gain enthalpy and it decreases down the group.

82. Electronegativity- Decreases down the group.
Colour- Halogens are coloured. Colour deepens down the group.

83. Oxidation States and Trends in Chemical Reactivity
All halogens exhibit -1 oxidation state.
Cl, Br and I exhibit +1, +3, +5 and +7 oxidation states.
F has absence of d- orbitals.
All halogens are highly reactive.
Reactivity decreases down the group.

84. Anomalous Behaviour of Fluorine
Small size, highest electronegativity, low F-F bond dissociation enthalpy and non- availability of d- orbital.
Reactions of fluorine are Exothermic.
HF is liquid while others are gases.

85. Reactivity towards Hydrogen
Form HX type.
The affinity for hydrogen decreases from fluorine to iodine.
Acidic strength increases from HF to HI.
Reducing character increases down the group.
Thermal stability decreases down the group.

86. Reactivity towards Oxygen
Halogens forms oxides with oxygen but most of them are unstable.
Fluorine forms two types of oxides- OF2(+1) and O2F2(+2).
Only O2F2 is thermally stable

87. Reactivity towards metals
Form Metal Halides.
Ionic character decreases from MF to MI.
The Halides in higher oxidation state will be more covalent than the one in lower oxidation state.

88. Reactivity of Halogens with other Halogens.
Form inter- Halogen Compounds like XX`, XX3`, XX5`, XX7` [ where X= larger size Halogen and X` is smaller size Halogen]

89. Chlorine Preparation –
MnO2 + 4 HCl MnCl2 + Cl2 + 2 H2O(heat)(conc.HCl)
4NaCl + MnO2 + 4 H2SO MnCl2 + 4 NaHSO4 + 2 H2O + Cl2
2KMnO4 +16 HCl KCl + 2 MnCl2 + 8H2O + 5Cl2

90. Manufacture of Cl2-
Deacon`s Process- By Oxidation
4 HCl + O Cl2+2H2O (CuCl2,723K)
2. Electrolytic Process- By electrolysis of brine (conc. NaCl solution). Cl2 is liberated at anode.

91. Properties of Chlorine
React with metals and non- metals to form chlorides.
2 Al + 3 Cl AlCl3
P4 + 6 Cl PCl3
It has great affinity for hydrogen.
H2 + Cl HCl
H2S + Cl HCl + S

92. Excess ammonia
8NH3 + 3Cl NH4Cl +N2
Excess chlorine
NH3 + 3 Cl NCl3 + 3HCl
Cold and dilute NaOH
2NaOH + Cl NaCl + NaOCl(hypochlorite) + H2O
Hot and conc. NaOH
6NaOH + 3Cl NaCl + NaClO3(chlorate) + 3H2O

93. With dry slaked lime
2Ca(OH)2 + 2Cl Ca(OCl)2+CaCl2+ 2H2O
Substitution Reaction-
CH4 + Cl CH3Cl(sunlight)
Oxidizes Fe2+ to Fe3+ and SO3- to SO42-
2FeSO4+H2SO4+Cl2 Fe2(SO4)3+2HCl

94. Powerful bleaching agent.
Cl2 + H2O HCl + [O]
Coloured subs. + [O] Colourless subs.

95. Hydrogen Chloride Preparation- Lab method-
NaOH+H2SO4(conc.) NaHSO4+HCl (420K)
NaHSO4 +NaCl Na2SO4+HCl (823K)
HCl can be dried by passing through conc. H2SO4.

96. Properties-
Ionizes as-
HCl + H2O H3O+ + Cl-
NH3 + HCl NH4Cl
3 conc. HCl +1 conc.HNO3 =Aqua Regia
It is used for dissolving platinum and gold.

97. Au + 4H+ + NO3- + 4Cl- AuCl4- + NO +2H2O
3 Pt + 16 H+ + 4 NO Cl PtCl NO +8 H2O
4. Na2CO3 + 2 HCl NaCl + H2O + CO2
NaHCO3 + HCl NaCl + H2O + CO2
Na2SO3 + 2HCl NaCl + H2O +SO2

98. Oxoacids of Halogens OxidationState/Chlorine/Bromine/Iodine
+1/HClO Hypochlorous acid/HBrO Hypobromous acid/HIO Hypoiodous acid
+3/HClO2 Chlorous acid / - / -
+5/HClO3 Chloric acid / HBrO3 Bromic acid/HIO3 Iodic acid
+7/HClO4 Perchloric acid/HBrO4 Perbromic acid/HIO4 Periodic acid 

99. Due to high electronegativity and small size , Fluorine forms only one oxoacid , HOF known as Fluoric (I) acid or Hypofluorous acid.

100. Oxoacids Chlorous Acid and Hypochlorous Acid

101. Oxoacids of Halogens- Perchloric Acid And Chloric Acid

102. Inter - Halogen Compounds
Preparation-
Equal volume
Cl2 + F ClF (437K)
Cl2 in excess
I2 + 3 Cl ICl3
Excess F2
Cl2 + 3 F ClF3 ( 573K )

103. Equimolar
I2 + Cl ICl
Diluted with water
Br2 + 3 F BrF3
Excess Fluorine
Br2 + 5 F BrF5

104. Properties-
Covalent compounds.
More reactive than the constituent halogen ( A-X bond is relatively weaker than X-X bond).
Diamagnetic in nature.
Good oxidising agent.
Melting and Boiling point increases with increase in difference in electronegativity.

105. Structure of Interhalogen Compounds
AX type:  In AX type, A is the central atom and X is the bonded atom.  In these type of compounds, both atoms share their electrons and form single covalent bond.
                                           A----X
                                                        

106.   In AX3 type, A is the central atom and undergoes sp3d1 hybridisation in first excited state and gives 5 hybrid orbitals.  Out of five, three are bonding electrons and two are lone pairs.  Since sp3d1 hybridisation, it should be in Trigonal bipyramidal shape, but due to the presence of two lone pairs, structure is distorted to "T"-- shape.
                                                        

107. AX5type: In AX5, A undergoes Sp3d2 hybridisation
AX5type:  In AX5, A undergoes Sp3d2 hybridisation.  The shape of the molecule is distorted octahedral or distorted square pyramidal.

108. AX7 type:  A undergoes sp3d3 hybridisation, and the shape of the molecule becomes pentagonal bipyramidal.

109. Group – 18 Elements Electronic configuration- ns2np6
Atomic Radii- Increases down the group.
Ionisation Enthalpy- Very high. Decreases down the group.
Electron gain enthalpy- Do not accept electrons and have large +ve values of electron gain enthalpy.
M.P. and B.P.- Low melting and boiling points.

110. Chemical Properties Least reactive. The inertness is due to:
Stable closed shell electronic configuration.
Exceptionally high ionisation enthalpy.
Very low electron affinities.

111. Xenon – Fluorine Compounds
Form 3 binary fluorides- XeF2 , XeF4 and XeF6.
Xe (excess) + F XeF2(673K,1 bar, Ni tube)
1:5 ratio
Xe +2F XeF4(873K,7 bar,Ni tube)
1:20 ratio
Xe +3F XeF6( 573K,60-70bar,Ni )

112. XeF4 + O2F XeF6 + O2
2 XeF2 +2H2O Xe + 4 HF + O2
XeF2 +PF [XeF]+[PF6]-
XeF4 +SbF [XeF3]+[SbF6]-
XeF6 + MF M+[XeF7]-(M=Na,K,Rb,Cs)

113. Structures