p block elements class 12

3. Some Exceptional Behavior
Considerable increase in covalent radius from nitrogen
to phosphorus but only a small increase from arsenic to
bismuth
This is due to completely filled d and f orbital's in
heavier members.
The ionization energy of group 15 elements is much
greater than the group 14elements in corresponding
period
This is due to extra stability of half filled p orbitals

4.  All elements are polyatomic
 Nitrogen is gas and all others are solid
 Metallic character increases down the group
 Except nitrogen all other elements show
allotropy
 Boiling point increases down the group
 Melting point increases up to As and then
decreases to Bi

5. Nitrogen differs from the rest of the members of
its group due to it's small size, high electro
negativity, high ionization enthalpy and non
availability of d-orbital’s.
It has an ability to form pπ-pπ bonds with itself
and hence it is inert at room temperature. Other
elements if its group are singly bonded.

6. • Towards Hydrogen (H2)
forming: EH3 (where E – N, P, As, Sb,
Bi)
• Towards Oxygen (O2)
forming E2O3 and E2O5 (where E – N, P,
As, Sb, Bi)
• Towards Metals

7. Preparations:
1. Labarotory preparation
NH4Cl(aq) + NaNO2(aq) → N2(g) + NaCl(aq) + 2 H2O (l)
2. thermal decompositon of ammonium
dichromate
(NH4)2Cr2O7(s) → Cr2O3(s) + N2(g) + 4 H2O(g)
3. Thermal decomposition of Na or Ba azide
2NaN3 → 2 Na + 3 N2

8. Colourless, odourless, tasteless, non toxic gas
Low molecular mass, low intermolecular
forces
Two stable isotopes
Low solubility in water , low freezing and
boiling point

9. ᧖ Combination with metals:
6 Li + N2 → 2 Li3N
᧖ Combines with hydrogen (H2):
(Haber’s process)
N2 + 3 H2 → 2 NH3 (ΔH = −92.4 kJ·mol−1)
᧖ Formatio of nitric acid (NO)
N2 + O2 → 2 NO

10. Manufacture of ammonia.
Provide an inert atmosphere in industries.
Used as a refrigerants.

11. Preparation:
o By Haber’s process
N2 + 3 H2 → 2 NH3 (ΔH = −92.4 kJ·mol−1)
o From ammonium salts
Ca(OH)2(s) + 2NH4Cl(s) → CaCl2(s) + 2H2O(l) + 2NH3(g)
o From decay of organic matter
2H2O + NH2CONH2 → (NH4)2 CO3 → 2 NH3 + 4H2O +
CO2

12. • Colourless with pungent odour
• Hydrogen bonding, thus has higher boiling
and melting point than expected on the basis
of molecular mass
• Structure- trigonal planar

13. Solubility in water
NH3 + H2O → NH4
+ + OH−
Reaction with acids
FeCl3 (aq) + 3NH4OH (aq) → Fe(OH)3 +
3NH4Cl
ZnSO4(aq) + 2 NH4OH(aq) → Zn(OH)2 +
(NH4)2(SO4)

14. • Production of various nitrogenous
fertilizers
• Manufacture of inorganic nitrogen
compounds
• Used as a refrigerants

15. Oxide Oxidation state characters
N2O nitrogen oxide
(nitrous oxide)
+1 Colourless and neutral
NO nitrogen monoxide
(nitric oxide)
+2 Colourless and neutral
N2O3 nitrogen trioxide +3 Blue , solid acidic
NO2 nitrogen dioxide +4 Brown, acidic
N2O4 nitrogen tetraoxide +4 Colourless and acidic
N2O5 nitrogen pentaoxide +5 Colourless and acidic

17. Preparation
Laboratory Preparation
It is prepared by heating KmnO3 or NaNO3 and
conc. H2SO4 in a glass resort.
NaNO3 + H2SO4 → NaHSO4 + HNO3
Industrial Preparation
On large scale it is prepared by Ostwald’s
process

18. Step 1 - Catalytic oxidation of NH3 by
atmospheric oxygen using Pt as catalyst at
500K and 9 bar pressure
4NH3 + 5O2 → 4NO (g) + 6H2O (g)
Step 2 – Nitric oxide combines with oxygen
2NO (g) + O2 → 2NO2 (g)
Step 3 – Nitrogen dioxide dissolves in water
3NO2(g) + H2O(l) → 2HNO3 (aq) + NO (g)

19. • Physical Properties
Commercially available nitric acid is an azeotrope with
water at a concentration of 68% HNO3, which is the
ordinary concentrated nitric acid of commerce. This
solution has a boiling temperature of 120.5 °C at
1 atm. Two solid hydrates are known; the
monohydrate (HNO3·H2O) and the trihydrate
(HNO3·3H2O).
• Nitric acid 70%
• Nitric acid of commercial interest usually consists of
the maximum boiling azeotrope of nitric acid and
water, which is approximately 68% HNO3, (approx. 15
molar). This is considered concentrated or technical
grade, while reagent grades are specified at 70%
HNO3. The density of concentrated nitric acid is
1.42 g/mL. An older density scale is occasionally seen,
with concentrated nitric acid specified as 42° Baumé.

20. 1. Reaction with elements less electropositive than
hydrogen
Concentrated HNO3
Cu + 4 HNO3 → Cu(NO3)2 + 2 NO2 + 2 H2O
Dilute HNO3
8 HNO3 + 3 Cu → 3 Cu(NO3)2 + 2 NO + 4 H2O

21. 2. Reaction with elements more electropositive
than hydrogen
Concentrated HNO3
4 Zn + 10 HNO3 (dilute) → 4 Zn(NO3)2 + N2O
+ 5 H2O
Dilute HNO3
Zn + 4 HNO3 (Conc) → Zn(NO3)2 + 2 NO2 + 2
H2O
3. Some metals donot dissolve in concentration
nitric acid because of formation of passive film
of oxide of surface.

22. This test is carried for checking the presence of
nitrate ion and it depends on the ability of Fe2+
nitrate to nitric oxide
Ferrous sulphate is added to aq. Solution and
sulphuric acid is added along the sides of the
test tube.
NaNO3 + FeSO4 +H2SO4 → [Fe(H2O)5 (NO)]SO4

23. Properties White red black
Colour White but yellow on
exposure
Dark red Black
State Waxy solid Brittle powder Crystalline
Stability Less stable More stable most stable
Chemical reactivity Very reactive Less reactive Least reactive

24. Preparation :
1. using calcium phosphide with water or dilute
HCl
Ca3P2 + 6HCl → + 2PH3
Ca3P2 + 6H2O→ + 2PH 2. Laboratory preparation3
P4 + 3NaOH + 3H2O (conc. And hot) →
PH3 + 3 NaH2PO2

25. • Colourless gas with rotten fish smell.
• Highly poisonous.
• It explodes in contact with traces of oxidising
agents like HNO3 .
• It is slightly soluble in water.
• It is weakly basic
PH3 + HBr → PH4Br

26. • Used in Homes signals
• Also used in smoke screens.

31. The elements of Group 16 have 6 electrons
in outermost shell and have ns2 np4 general
electronic configuration
• Oxygen: 1s2 2s2 2p4
• Sulphur: 1s2 2s2p6 3s2p4
• Selenium: 1s2 2s2p6 3s2p6d10 4s2p4
• Tellurium: 1s2 2s2p6 3s2p6d10 4s2p6d10 5s2p4
• Polonium:
1s2 2s2p6 3s2p6d10 4s2p6d10f14 5s2p6d10 6s2p4

33. Due to increase in number of shells, atomic and ionic radii
increases from top to bottom in group. The size of oxygen atom
is exceptionally small.
Ionisation enthalpy of elements of group 16 is
lower than group 15 due to half filled p-orbitals in group
15 which are more stable. However, ionization enthalpy
decreases down the group.

34. Oxygen has less negative electron gain
enthalpy than S because of small size of O.
From S to Po electron gain enthalpy becomes less negative to Po
because of increase in atomic size.
Next to fluorine, oxygen has the highest electronegativity
value amongst the elements. With in the group,
electronegativity decreases with an increase in atomic
number. This implies that metallic character increases from
oxygen to polonium.

35. • Oxygen and sulphur are non-metals, selenium
and tellurium are metalloids, whereas polonium
is a metal.
• All the elements exhibit allotropy.
• The melting point and boiling point increase with
an increase in atomic number down the group.
• The large difference in melting and boiling point
of oxygen (diatomic) and sulphur (polyatomic) is
due to difference in atomicity.

36. Oxidation states and trends in chemical reactivity
1) They show -2, +2, +4, +6 oxidation states. Oxygen does not show
+6 oxidation state due to absence of d –orbitals. Po does not show
+6 oxidation state due to inert pair effect.
2) The stability of -2 oxidation state decreases down the group due to
increase in atomic size and decrease in electronegativity.
3) Oxygen shows -2 oxidation state in general except in OF2 and
O2F2.
4) The stability of +6 oxidation state decreases and +4 oxidation state
increases due to inert pair effect.

37. Anomalous behaviour of oxygen
• The anomalous behaviour of oxygen is due to its
small size and high electronegativity.
• The absence of d orbitals in oxygen limits its
covalence to four and in practice, rarely exceeds
two. On the other hand, in case if other elements
of the group, the valence shells can be expanded
and covalence exceeds 4.

38. Reactivity with hydrogen
All group 15 elements from trihydrides, MH3. Hybridisation - sp3
The stability of hydrides decrease down the group due to decrease in
bond dissociation energy down the group.
NH3 > PH3 > AsH3 > SbH3 > BiH3
Boiling point:
PH3 < AsH3 < NH3 < SbH3 < BiH3
Boiling point increases with increase in size due to increase in van der
Waals forces. Boiling point of NH3 is more because of hydrogen bonding.
Bond angle:
NH3 (107.8°) > PH3 (99.5°) > AsH3 (91.8°) ≈ SbH3(91.3°) > BiH3 (90°)
Electronegativity of N is highest. Therefore, the lone pairs will be
towards nitrogen and hence more repulsion between bond pairs.
Therefore bond angle is the highest. After nitrogen, the
electronegativity decreases down the group.

39. All group 15 elements from trioxides (M2O3) and pentaoxides
(M2O5).
Acidic character of oxides decreases and basicity increases down
the group. This is because the size of nitrogen is very small. It has
a strong positive field in a very small area. Therefore, it attracts
the electrons of water’s O-H bond to itself and release H+ ions
easily.
As we move down the group, the atomic size increases. Hence,
the acidic character of oxides decreases and basicity increases as
we move down the group.

40. Group 15 elements form trihalides and pentahalides.
Trihalides – covalent compounds and become ionic down the group.
Sp3 hybridisation , pyramidal shape
Pentahalides - Sp3d hybridisation, TBP shape
They are lewis acids because of the presence of vacant d – orbitals.
PCl5 + Cl-→ [PCl6]-
PCl5 is ionic in solid state and exist as [PCl4]+ [PCl6]-
In PCl5, there are three equatorial bonds and two axial bonds. The
axial bonds are longer than equatorial bonds because of greater
repulsion from equatorial bonds.

42. It can be obtained in the laboratory by the following ways:-
 By heating oxygen containing
salts such as chlorates, nitrates, & permanganates.
2KClO3 -----> 2KCl + 3O2
By the thermal decomposition of the oxides of metals
low in the electrochemical series & higher oxides of some
metals.
2Ag2O ------> 4Ag+ O2
2Pb3O4 --------> 6PbO + O2
Hydrogenperoxide is readily decomposed into water and
dioxygen by catalysts such as finely divided metals and
manganese dioxide
2H2O2----->2H2O + O2
On large scale it can be prepared from water or air.

43.  Dioxygen is a colourless & odourless gas.
Its solubility in water is to the extent of
3.08cubic cm in 100cubic cm water at 293K which
is just sufficient for the vital support of marine
and aquatic life.
It liquefies at 90K and freezes at 55K.
It directly reacts with nearly all metal & non-metals
& some noble gases

44. Its importance in normal respiration.
Combustion processes.
It is used in oxyacetylene welding.
In the manufacture of many metals ,
particularly steel.
Oxygen cylinders are widely used in hospitals
, high altitudes flying and in mountaineering.
For combustion of fuel.

45. Ozone is an allotropic form of oxygen. It is too
reactive to remain for long in atmosphere at sea
level at the height of about 20km.

46. When a slow dry stream of oxygen is passed
through a silent electrical discharge
3O2------> 2O3 = +142 kJ/mol
it is an endothermic process.
If concentration of ozone greater than 10% are
required , a battery of ozonisers can be used & pure
ozone (b.p.385K) can be condensed in a vessel
surrounded by liquid oxygen.

47. Pure ozone is a pale blue gas , dark blue liquid & violet-black
solid.
Its small concentration is harmless but if concentration
rises above 100ppm breathing becomes uncomfortable
resulting in headache and nausea.
It is thermodynamically unstable with respect to oxygen.
High concentration of ozone can be dangerously
explosive.
It acts as a powerful oxidising agent.
PbS + 4O3------->PbSO4 + 4O2

48. It is used as a germicide, disinfectent, for
sterilising water.
It is also used for bleaching oils , ivory, flour ,
starch,etc.
It acts as an oxidising agent in the manufacture
of potassium permanganate.

49. Sulphur forms numerous allotropes of which the yellow rhombic( alpha-sulphur)
& monoclinic ( beta-sulphur) forms are the most important.

50. It is yellow in colour.
Its m.p. 385.8 K
Specific gravity 2.06.
Its crystals are formed on evaporating the
solution of roll sulphur in CS2 .
It is insoluble in water but dissolves to some extent in
benzene, alcohol, & ether
It is readily soluble in CS2.

51. Its melting point is 393K.
Specific gravity is1.98
It is soluble in CS2
It is colourless.
It has needle shaped crystals.
BOTH THE MOLECULES HAVE S8 MOLECULES. THE S8
RING IN BOTH THE FORMS IS PUCKERED AND HAS A
CROWN SHAPE.

53. Sulphur dioxide is formed together with a little (6-
8%) sulphur trioxide when sulphur is burnt in air or
oxygen.
S + O2 ------> SO2
In laboratory it is readily generated by treating a sulphite
with dil. Sulphuric acid
It is produced as a bi-product of the roasting of
sulphide ores.
4FeS2 + 11O2------>2Fe2O3 + 8SO2
THE GAS AFTER DRYING IS LIQUEFIED UNDER PRESSURE
AND STORED IN STEEL CYLINDERS.

54. Sulphur dioxide is a colourless gas with pungent smell.
It is highly soluble in water.
It liquefies at room temperature under 2atm pressure.
It boils at 263K.
It reacts readily with NaOH solution, forming sodium
sulphite.
2NaOH + SO2------> Na2SO3 + H2O
In its reaction with water and alkalies , the behaviour of
sulphur dioxide is very similar to that of carbon dioxide.
WHEN MOIST, SULPHUR DIOXIDE BEHAVES AS A REDUCING
AGENT
2Fe(III) + SO2 + 2H2O-----> 2Fe(II) +SO4(-2) + 4H+

55. It is used in refining petroleum and
sugar.
It is used in bleaching wool and silk.
It is used as anti –chlor , disinfectant
and preservative.
Liq. SO2 is used as a solvent to dissolve a
number of organic & inorganic chemicals.

57. IT IS ONE OF THE MOST
IMPORTANT INDUSTRIAL
CHEMICALS WORLDWIDE ,
THEREFORE IT IS CALLED AS KING
OF CHEMICALS……….!!

58. Industrially Sulphuric acid is manufactured by CONTACT
PROCESS….
The reaction is exothermic , reversible and the forward
reaction leads to decrease in volume.
In practice , the plant is operated at a pressure of 2 bar &
at temperature 720K.
Dilution of oleum with water gives H2SO4.
THE SULPHURIC ACID OBTAINED BY CONTACT PROCESS IS 96-
97% PURE….!!!

59. It is colourless liquid , dense , oily liquid.
It has specific gravity of 1.84 at 298K.
It freezes at 283K & boils at 611K.
Concentrated sulphuric acid is a dehydrating
agent.
C12H22O11 + H2SO4------> 12C + 11H2O
Hot concentrated sulphuric acid is a moderately
strong oxidising agent.
Cu + 2H2SO4(conc.)----> CuSO4 + SO2 + 2H2O
3S + 2H2SO4(conc.)-------> 3SO2 + 2H2O

60. It is used in petroleum refining.
Manufacture of pigments, paints &
dyestuff intermediates.
In detergent industry.
Storage batteries.
As a laboratory reagent.

62. 1. Outermost configuration is ns2np5.
2. They have very high ionization enthalpy due to very small atomic
size. The I.E. decreases as we go down the group.
3.They have maximum negative electron gain enthalpy in the
corresponding periods . This is because after gaining an electron
they attain the stable noble gas configuration.
However, the negative electron gain enthalpy of fluorine is less than
that of chlorine. It is due to small size of fluorine atom. As a result,
there are strong interelectronic repulsions in the relatively small 2p
orbitals of fluorine and thus, the incoming electron does not
experience much attraction.
4. All halogens are coloured. This is due to absorption of radiations in
visible region which results in the excitation of outer electrons to
higher energy level. For example, F2, has yellow, Cl2 , greenish
yellow, Br2, red and I2, violet colour

63. • 5.The enthalpy of dissociation of F2 is less
compared to that of Cl2
• A reason for this anomaly is the relatively
large electron-electron repulsion among
the lone pairs in F2 molecule due to
smaller size of F-atom where they are
much closer to each other than in case of
Cl2.
• Cl – Cl > Br – Br > I – I.
• Cl2> Br2>F2>I2

64. Although electron gain enthalpy of fluorine
is less negative as compared to chlorine,
fluorine is a stronger oxidising agent than
chlorine. Why?

65. Answer :- It is due to
(i) low enthalpy of dissociation of F-F bond
(ii) high hydration enthalpy of F–

66. 1.Fluorine shows only -1 state . due to its high electronegativity.
2. The fluorine atom has no d orbitals in its valence shell and
therefore cannot expand its octet.
3. Others show -1 , +1, +3, +5 and +7.
4. The chemical reactivity decreases down the group, b’coz the
electronegativity decreases.
5. The Oxidising tendency also decreases down the group whereas
the reducing tendency increases.
There is a regular decrease in the first ionization energy as we go
down this column. As a result, there is a regular decrease in the
oxidizing strength of the halogens from fluorine to iodine. F2 >
Cl2 > Br2 > I2
F2 > Cl2 > Br2 > I2
oxidizing strength

67. 1.The order of boiling point is:-
HF > HI> HBr> HCl
HI has exceptionally higher bpt since it forms extensive
H-bonding.
The remaining hydrides follow increase in bpt with
increase in molar mass.
1. Acidic strength of these acids increases in the order:
.
Reason:- BDE decreases.
H–F > H–Cl > H–Br > H–I
2. The stability of these halides decreases down the
group
H–F > H–Cl > H–Br > H–I.

68. 1. Fluorine forms two oxides OF2 and O2F2.
However, only OF2 is thermally stable at
298 K.
These oxides are essentially oxygen fluorides
because of the higher electronegativity of
fluorine than oxygen.

69. Fluorine exhibits only –1 oxidation state
whereas other halogens exhibit + 1, + 3, +
5 and + 7 oxidation states also. Explain.

70. Answer :-
Fluorine is the most electronegative element
and cannot exhibit any positive oxidation
state. Other halogens have d orbitals and
therefore, can expand their octets and
show + 1, + 3, + 5 and + 7 oxidation
states also.

71. Preparation
Laboratory preparation
1. By heating manganese dioxide with concentrates HCl.
MnO2 + 4 HCl ==> MnCl2 + 2 H2O + Cl2
2. By the action of HCl on potassium permanganate.
2 KMnO4 + 16 HCl ==> 2 MnCl2 + 2 KCl + 8 H2O + 5 Cl2
Manufacture of chlorine
1. Deacon’s process
By the oxidation of HCl gas by atmospheric oxyden in presence of CuCl2
723 k.
2. Electrolytic process
Chlorine is manufactured industrially as a by-product in the manufacture of
Caustic Soda by the electrolysis of brine.
2 NaCl + 2 H2O ==> Cl2 + H2 + 2 NaOH

72. Chlorine is
• a highly toxic greenish yellow gas,
• has a pungent odour, and
• fumes in moist air.
Reaction of chlorine with metals
2Al + 3Cl2 ==> 2 AlCl3
2Na + Cl2 ==> 2NaCl
Reaction with Hydrogen
H2 + Cl2 ==> 2HCl
H2S + Cl2 ==> 2HCl + S

73. With excess of ammonia, Cl gives nitrogen and ammonium chloride whereas
with excess of chlorine it gives nitrogen trichloride (explosive)
8NH3 + 3Cl2 ==> 6NH4Cl + N2
NH3 + 3Cl2 ==> NCl3 + 3HCl
With cold and dilute alkalies chlorine produces a mixture of chloride and
hypochlorite but with hot and concentrated alkalies it gives chloride and
chlorate
2NaOH + Cl2 ==> NaCl + NaOCl + 3H2O
(Cold and Dilute)
6NaOH + 3Cl2 ==> 5NaCl + NaClO3 + 3H2O
Preparation of Bleaching Powder
2 Cl2 + 2 Ca(OH)2 → Ca(OCl)2 + CaCl2 + 2 H2O

74. Bleaching action of Chlorine
Cl water on standing loses its yellow colour due to formation of
HCl and HOCl
Bleaching action is due to oxidation.
Cl2 + H2O → 2HCl + O Reaction showing bleaching action of Cl

75. • for the manufacture of bleaching powder and liquid bleaches,
• to bleach fabrics (e.g. linen and cotton), wood pulp and paper,
• in the manufacture of a wide range of chloro-organic solvents, including
Methylene Chloride, CH2Cl2, Chloroform, CHCl3, Carbon Tetrachloride, CCl4,
• in the manufacture of a number of important inorganic chemicals, including
Sulphur Chloride, S2Cl2, Thionyl Chloride, SOCl2, Phosgene (i.e. Carbonyl
Chloride), COCl2, and inorganic Chlorates, (e.g. Sodium Chlorate, NaClO3),
• as a disinfectant used to kill bacteria in the preparation of drinking water.
• Chlorine is also important in the manufacture of paints, aerosol propellants and
plastics.
• for the extraction of Gold from its ores,

76. PREPARATION
Laboratory methods
1. NaCl + H2SO4 → NaHSO4 + HCl
2. NaCl + NaHSO4 → HCl + Na2SO4

77. 1. It is a colourless and pungent smelling gas.
2. It is easily liquefied and freezes to a white crystalline solid.
3. It is extremely soluble in water.
4. It reacts with NH3 and gives white fumes of NH4Cl.
NH3 + HCl → NH4Cl
5. When 3 parts of concentrated HCl and one part of
concentrated HNO3 are mixed aqua regia is formed which is used
to dissolve noble metals.
6. HCl decomposes salts of weaker acids.
Na2CO3 + 2HCl 2NaCl + H2O + CO2

78. Fluorine forms only one oxoacid HOF (Fluoric (I)
acid or hypofluorous
acid) due to high electronegativity.
Acid strength: HOCl < HClO2 < HClO3 < HClO4
Reason: HClO4 → H++ ClO4
-
Acid strength: HOF > HOCl > HOBr > HOI
This is because Fluorine is most electronegative.

79. Binary compounds of two different halogen atoms of general formula
X X’n are called interhalogen compounds where n = 1, 3, 5, or 7
These are covalent compounds.
All these are covalent compounds.
Interhalogen compounds are more reactive than halogens because XX’ is a
more polar bond than X-X bond.
All are diamagnetic.
Their melting point is little higher than halogens.
XX’ (CIF, BrF, BrCl, ICl, IBr, IF) (Linear shape)
XX’3 (CIF3, BrF3, IF3, ICl3) (Bent T- shape)
XX’5 – CIF5, BrF5, IF5, (square pyramidal shape)
XX’7 – IF7 (Pentagonal bipyramidal shape)

81. • 1.Group 18 consists of six elements: helium, neon,
argon, krypton, xenon and radon.
• 2.Electronic configuration is ns2np6
• 3.They have very high ionization enthalpy due to
stable configuration.
• 4..All these are gases and chemically unreactive.
They form very few compounds. Because of this
they are termed noble gases.
• 5.All the noble gases except radon occur in the
atmosphere. Their atmospheric abundance in dry
air is ~ 1% by volume of which argon is the major
constituent.

82. Helium and sometimes neon are found in minerals of
radioactive origin e.g., pitchblende, monazite, cleveite.
The main commercial source of helium is natural gas.
Xenon and radon are the rarest elements of the group.
6.They have very low melting and boiling points
because the interatomic interaction in these elements
is weak dispersion forces. Helium has the lowest
boiling point (4.2 K) of any known substance. It has an
unusual property of diffusing through most commonly
used laboratory materials such as rubber, glass or
plastics.

83. • 7.In general, noble gases are least reactive.
Their inertness to chemical reactivity is
attributed to the following reasons:
• (i) The noble gases except helium (1s2)
have completely filled ns2np6 electronic
configuration in their valence shell.
• (ii) They have high ionisation enthalpy and
more positive electron gain enthalpy.(since
they do not have any tendency to accept
an additional electron as their outermost
orbits are completely filled)

84. 8.Xenon has lower ionisation enthalpy,
because of large size of xenon.
9. Xenon reacts with only oxygen and
fluorine atoms because these two elements
are of very high polarizing capacity.
10.The fact that the first IE of xenon is
comparable with that of molecular oxygen
prompted Neil Bartlett to study the
chemistry of xenon compounds.(he
understood this from the compound O2PtF6,
which he prepared )
The first compound of xenon prepared was
XePtF6

85. • Why are the elements of Group 18 known as
noble gases ?

86. The elements present in Group 18 have their
valence shell orbitals completely filled and,
therefore, react with a few elements only
under certain conditions. Therefore, they are
known as noble gases.

87. • Xenon forms three binary fluorides, XeF2,
XeF2 and XeF6 by the direct reaction of
elements under appropriate experimental
conditions.
• XeF6 can also be prepared by the interaction
of XeF4 and O2F2 at 143K.
• XeF4 + O2 F2
 XeF6 + O2

88. 1.XeF2 is hydrolysed to give Xe, HF and O2.
2XeF2 (s) + 2H2O(l)2Xe (g) + 4 HF(aq) +
O2(g)
2.Xenon fluorides react with fluoride ion
acceptors to form cationic species and
fluoride ion donors to form fluoroanions.

89. Hydrolysis of XeF4 and XeF6 with water gives
XeO3.
6XeF4 + 12 H2O 4Xe + 2XeO3 + 24 HF + 3
O2
XeF6 + 3 H2O XeO3 + 6 HF

92. 1.Helium is a non-inflammable and light gas. Hence, it
is used in filling balloons for meteorological
observations. It is also used in gas-cooled nuclear
reactors. Liquid helium (b.p. 4.2 K) finds use as
cryogenic agent for carrying out various experiments
at low temperatures. It is used to produce and
sustain powerful superconducting magnets which
form an essential part of modern NMR
spectrometers and Magnetic Resonance Imaging
(MRI) systems for clinical diagnosis. It is used as a
diluent for oxygen in modern diving apparatus
because of its very low solubility in blood.
.

93. 2.Neon is used in discharge tubes and
fluorescent bulbs for advertisement display
purposes.
3.Argon is used mainly to provide an inert
atmosphere in high temperature metallurgical
processes (arc welding of metals or alloys) and
for filling electric bulbs.
It is also used in the laboratory for handling
substances that are air-sensitive.
4.Xenon and Krypton are used in light bulbs
designed for special purposes

95. Compiled and
presented by
Kush Sehgal r.no 29
Lakshay Thakur r.no 30
Mamik Dutta r.no 31
Manisha Dash r.no 32
Mayank Kashyap r.no