Chapter 12 s & p Block Elements part 3 by TEACHING CARE Online coaching and tuition classes

November 13, 2021 admin Off Uncategorized

Chemical Fertilizers : The chemical substances which are added to the soil to keep up the fertility of soil are called

Types of fertilizers : Chemical fertilizers are mainly of four types,

Nitrogenous fertilizers : g. Ammonium sulphate

NH ₂CONH ₂ etc.

(NH 4 )2 SO4 ,

Calcium cyanamide

CaCN 2 ,

Urea

Phosphatic fertilizers : g. Ca(H ₂ PO₄ )₂ .H ₂O (Triple super phosphate), Phosphatic slag etc.
Potash fertilizers : g. Potassium nitrate (KNO₃ ), Potassium sulphate (K₂SO₄ )etc.
Mixed fertilizers : These are made by mixing two or more fertilizers in suitable e.g. NPK

(contains nitrogen, phosphorus and potassium).

NPK is formed by mixing ammonium phosphate, super phosphate and some potassium salts.

Oxygen is the first member of group 16 or VIA of the periodic table. It consists of five elements Oxygen (O), sulphur (S), selenium (Se), tellurium (Te) and polonium (Po). These (except polonium) are the ore forming elements and thus called chalcogens.

Electronic configuration

Elements Electronic configuration (

ns ² np⁴

)

8 O 1s ² ,2s ² 2p⁴ or [He] 2s ² 2p ⁴

16 S 1s ² ,2s ² 2p⁶ ,3s ² 3p⁴ or [Ne] 3s ² 3p ⁴

34 Se 1s ² ,2s ² 2p⁶ ,3s ² 3p⁶ 3d¹⁰ ,4s ² 4 p⁴ or [Ar] 3d¹⁰ 4s ² 4 p ⁴

52 Te 1s ² ,2s ² 2p⁶ ,3s ² 3p⁶ 3d¹⁰ ,4s ² 4 p⁶ 4d¹⁰ ,5s ² 5p ⁴ or [Kr] 4d¹⁰ 5s ² 5p ⁴

84 Po 1s ² ,2s ² 2p⁶ ,3s ² 3p⁶ 3d¹⁰ ,4s ² 4 p⁶ 4d¹⁰ 4 f ¹⁴ ,5s ² 5p⁶ 5d¹⁰ ,6s ² 6p ⁴ or [Xe] 4 f ¹⁴ 5d¹⁰ 6s ² 6p ⁴

Physical properties

Physical state : Oxygen is gas while all other are
Atomic radii : Down the group atomic radii increases because the increases in the number of inner shells overweighs the increase in nuclear
Ionisaion energy : Down the group the ionisation energy decrease due to increase in their atomic radii and shielding
Electronegativity : Down the group electronegativity decreases due to increase in atomic
Electron affinity : Element of this group have high electron affinity electron affinity decreases down the
Non – metallic and metallic character : These have very little metallic character because of their higher ionisation
Nature of bonding : Compound of oxygen with non metals are predominantly S, Se, and Te because of low electronegativities show more covalent character.
Melting and boiling points : The pt. and b. pt increases on moving down the group.
Catenation : Oxygen has some but sulphur has greater tendency for

H – O – O – H, (H2O2)

H – S – S – H, (H2S2)

H – S – S – S – H, (H2S3 )

H – S – S – S – S – H

(H2S4 )

(10) Allotropy

*Oxygen –*		O2 *and* O3
Sulphur –		Rhombic , monoclinic, plastic sulphur
Selenium–		Red (non-metallic) grey (metallic)
Tellurium	–	Non-metallic and metallic (more
		stable)
Polonium	–	a and b (both metallic)

Oxidation states : Oxygen ® – 2 and –1 oxidation These element shows +2 ,+4 and +6 oxidation state.

Chemical properties

(1) Hydrides :

H 2 O, H 2 S, H 2 Se, H 2 Te

and

H 2 Po, H 2 O – colourless and odourless.

H 2 S, H 2 Se, H 2 Te

and

H ₂ Po – colourless, unpleasant smell.

Increasing order of reducing power of hydrides :

H 2 O < H 2 S < H 2 Se < H 2 Te

Increasing order of bond angles in hydrides :

H 2 Te < H 2 Se < H 2 S < H 2 O

The order of stability of hydrides :

H 2 O > H 2 S > H 2 Se > H 2 Te

The order of increasing acidic nature of hydrides :

H 2 O < H 2 S < H 2 Se < H 2 Te

Oxides : These elements form monoxides (MO), dioxides ( MO₂ ) increasing order of acidic nature of oxides is TeO3 < SeO3 < SO3 .

(3) Oxyacids :

H 2 SO3 , H 2 SO4 , H 2 S2 O3 , H 2 SO5 , H 2 S2 O8 , H 2 S2 O7 , H 2 S2 O6

Halides : Oxygen : OF2 , Cl2 O, Br2 O

Sulphur : S2 F2 , S2 Cl2 , SF2 , SCl2 , SBr2 , SF4 , SCl4

and

SF6

Selenium and tellurium : SeF6 and TeF6

Oxygen and its compounds

Oxygen is the most abundant element in the earth crust (46.5%). It was discovered by Karl Scheele and

Joseph Priestley. It occurs in three isotopic forms :

8 O16

( Abundance: 99.76%)

Out of the three isotopes, ₈ O¹⁸

8 O17

( Abundance: 0.037%)

is radioactive.

8 O18

( Abundance: 0.204%)

Occurrence : In free state, it occurs in air and constitutes 21% by volume of
Preparation of Dioxygen : Oxygen is prepared by the following
- By the decomposition of oxygen rich compounds : g.

2KNO3 ¾¾^H^e¾^a^t ® 2KNO2 + O2 ; 2KClO3 ¾¾^H^e¾^a^t ® 2KCl + 3O2

Pot. Nitrate

Pot. Chlorate

MnO2

By heating dioxides, Peroxides and higher oxides : g.

2Ag 2 O ¾¾^H^e¾^a^t ® 4 Ag + O2 ;

Silver oxide

3MnO2

Manganese dioxide

¾¾^H^e¾^a^t ® Mn3 O4 + O2 ;

2BaO2

Barium peroxide

¾¾He¾at ®

2BaO

Barium oxide

Laboratory Method : In the laboratory, O₂ is prepared by thermal decomposition of potassium

2KClO3 ¾¾⁴²⁰¾^K ® 2KCl + 3O2

MnO2

Note : ® In the absence of

MnO2

catalyst, the decomposition takes place at 670-720 K. Therefore, MnO2

acts as a catalyst and also lowers the temperature for the decomposition of

KClO3 .

O2can also be prepared by the action of water on sodium peroxide as, 2Na2 O2 + 2H 2 O ® 4 NaOH + O2 .
Industrial preparation : The main sources for the industrial preparation of dioxygen are air and

From air :

O₂ is prepared by fractional distillation of air. During this process,

N ₂ with less boiling point

(78 K) distills as vapour while O₂ with higher boiling point (90 K) remains in the liquid state and can be separated.

From water :

O₂ can also be obtained by the electrolysis of water containing a small amount of acid or

alkali, 2H ₂O

Electrolysis

2H ₂ (g) + O₂ (g) .

Physical properties of O₂ : It is a colourless, tasteless and odourless It is slightly soluble in water and its solubility is about 30 cm³ per litre of water at 298 K.

Physical properties of atomic and molecular oxygen

Chemical properties of O₂ : It does not burn itself but helps in burning. It is quite stable in nature and its bond dissociation energy is very Therefore, it is not very reactive as such, O2 ® O + O .

Therefore, dioxygen reacts at higher temperatures. However, once the reaction starts, it proceeds of its own. This is because the chemical reactions of dioxygen are exothermic and the heat produced during the reaction is sufficient to sustain the reactions.

Action with litmus : Like dihydrogen, it is also neutral and has no action on blue or red
Reaction with metals : Active metals like Na, Ca react at room to form their respective oxides.

4 Na + O2 ® 2Na2 O ; 2Ca + O2 ® 2CaO

It reacts with Fe, Al, Cu etc. metals at high temperature

4 Al + 3O2 ® 2Al 2 O3 ; 4 Fe + 3O2 ® 2Fe2 O3

Action with Non-metals : It form

1073 K

2H 2 + O2 ¾¾Elec¾tric¾disc¾har¾ge ® 2H 2 O ;

N 2 + O2 ¾¾³²⁷¾³¾^K ®

2NO

Nitric oxide

S + O2 ¾¾^H^e¾^a^t ® SO2 ; C + O2 ¾¾^H^e¾^a^t ® CO2

Reaction with compounds : Dioxygen is an oxidising agent and it oxidises many compounds under

700 K

specific conditions. e.g. 4 HCl + O2 ¾¾CuC¾l ® 2H 2 O + 2Cl2 ;

4 NH 3 + 5O2 ¾¾¹⁰⁷¾³¾^K ® 4 NO + 6H 2 O

CS2 + 3O2 ¾¾^H^e¾^a^t ® CO2 + 2SO2 ;

(5) Uses of dioxygen

CH4 + 2O2 ® CO2 + 2H 2 O

It is used in the oxy-hydrogen or oxy-acetylene torches which are used for welding and cutting of
It is used as an oxidising and bleaching agent,

Liquid O₂

is used as rocket fuel.

It is used in metallurgical processes to remove the impurities of metals by

(6) Compounds of Oxygen

Oxides : A binary compound of oxygen with another element is called oxide. On the basis of acid-base characteristics, the oxides may be classified into the following four types,

Basic oxides : Alkali, alkaline earth and transition metals form basic oxides –

Na 2 O, MgO, Fe2 O3

etc. their

relative basic character decreases in the order : alkali metal oxides>alkaline earth metal oxides>transition metal oxides.

Acidic oxides : Non–metal oxides are generally acidic – CO2 , SO2 , SO3 , NO2 , N 2 O5 , P4 O10 , Cl2 O7

etc.

Amphoteric oxides :

Al2 O3 , SnO2 etc.

Neutral oxides :

H ₂ O, CO, N ₂O, NO

etc.

Trends of oxides in the periodic Table : On moving from left to the right in periodic table, the nature of the oxides change from basic to amphoteric and then to acidic. For example, the oxides of third period has the following behaviour,

Na2O

strongly basic

MgO basic

Al2 O3

amphoteric

SiO₂ weakly

acidic

P4 O10 acidic

SO₂ strongly

acidic

Cl₂O₇ very

strongly acidic

However, on moving down a group, acidic character of the oxides decreases. For example in the third group, the acidic character of oxides decreases as:

B₂O₃ acidic

Al₂O₃ amphoteric

Ga₂O₃ (weakly basic)

In2 O3 , Tl 2 O3 basic

On the basis of oxygen content the oxides may be classified into the following types,

Normal oxides : These contain oxygen atoms according to the normal oxidation number i.e. – 2. For example, MgO, H 2 O, CaO, Li2 O, Al 2 O3 etc.

Polyoxides : These contain oxygens atoms more than permitted by the normal valency. Therefore, these contain oxygen atoms in oxidation state different than –2.

Peroxides : These contains

H 2 O2 , Na2 O2 , BaO2 , PbO2 etc.

Superoxides : These contains

KO₂ , PbO₂ , etc.

²^– ion having oxidation number of oxygen as –1. For example,

ion having oxidation number of oxygen as –1/2. For example,

Suboxides : These oxides contain less oxygen than expected from the normal valency. For example,

N 2 O.

Mixed oxides : These oxides are made up of two simple oxides. For example, red lead

Pb3 O4 (2PbO2 + PbO2 ),

Mn₃O₄ (MnO₂ + 2MnO).

magnetic oxide of iron,

Fe3 O4 (FeO + Fe2 O3 )

and mixed oxide of manganese,

Ozone or trioxygen

Ozone is an allotrope of oxygen. It is present in the upper atmosphere, where it is formed by the action of

V. radiations on O₂ ,

3O2 ¾¾U.V¾.rad¾iatio¾n ® 2O3 .

Ozone

O₃ protects us from the harmful U. V. radiations which causes skin cancer. Now a days, ozone layer in the stratrosphere is depleting due to NO released by supersonic aircrafts and chlorofluoro carbons (CFC’S) i.e. freon which is increasingly being used in aerosols and as a refrigerant.

Preparation : Ozone is prepared by passing silent electric discharge through pure, cold and dry oxygen in a specially designed apparatus called ozoniser. The formation of ozone from oxygen is an endothermic reaction.

Silent electric

3O2 2O3

discharge

DH = +285.4 kJ

Ozone is prepared in the laboratory by the following two types of ozonisers,

(a) Siemen’s ozoniser, (b) Brodie’s ozoniser

For the better yield of ozone : (a) Only pure and dry oxygen should be used. (b) The ozoniser must be

perfectly dry. (c) A fairly low temperature sparkless.

(» 273 K)must be maintained. (d) The electric discharge must be

Physical properties : Ozone is a light blue coloured gas, having pungent odour. It is heavier than air. Its vapour density is 24. It is slightly soluble in water.

Chemical properties : The important chemical properties of ozone are discussed below,

Decomposition : Pure ozone decomposes on heating above 475 K to from O₂

gas.

2O 3 ¾¾⁴⁷⁵¾^K ® 3O2

DH = -285.4 kJ

Oxidising agent : Ozone is one of the most powerful oxidising agent with the liberation of In fact, ozone is a stronger oxidising agent than molecular oxygen because ozone has higher energy content and

decomposes to give atomic oxygen as: O3 ® O2 +

Atomic oxygen

Therefore, ozone oxidises a number of non-metals and other reducing agents. e.g.

2Ag+ O₃ ®

Metal

Ag 2 O + O2 ;

Silver oxide

Non–metal

+ 3O3 ® SO3 + 3O2 ;

PbS

Compound

+ 4O3 ® PbSO4 + 4O2

Mercury is oxidised to mercurous oxide,

2Hg + O₃ ®

Hg 2 O

Mercurous oxide

During this reaction mercury loses its meniscus and starts sticking to the sides of the glass. This is known as tailing of mercury. Mercurous oxide formed in this reaction dissolves in mercury and starts sticking to the glass surface.

Bleaching agent : Due to the oxidising action of ozone, it acts as a mild bleaching agent as well as a sterilizing It acts as a bleaching agent for vegetable colouring matter.

Vegetable

colouring

matter + O3 ® Oxidised coloured matter+ O2

(Colourless)

For example, ozone bleaches indigo, ivory, litmus, delicate fabrics etc.

Formation of ozonides : Ozone reacts with alkenes in the presence of CCl₄ to form an e.g.

CH 2 = CH 2 + O3 ¾¾^C^C¾^l4 ® H 2 C

Ethylene

CH 2

O O

Ethylene ozonide

Structure of O₃

: The structure of

O₃ molecule is angular as shown in

fig. The O – O – O bond angle is 116.8° and O – O bond length is 128 pm.

Uses of ozone

O₃ is used for disinfecting water for drinking purposes because ozone has germicidal
It is used for purifying air of crowded places such as cinemas, under ground railway, auditoriums, tunnels, mines

It is used in industry for the manufacture of

KMnO₄ , artificial silk, synthetic camphor etc.

Sulphur and its compounds

Sulphur is the second member of oxygen family and belongs to group-16 (VI A) of the periodic table.

Occurrence : Sulphur occurs in the earth’s crust to the extent of 05%. It occurs in the free state as well as in combined state. Sulphure occurs mainly as sulphides and sulphates. eg.

*Sulphide Ores Sulphate Ores*
Iron pyrites (fool’s gold)			– FeS2	Gypsum	– CaSO4 .2H 2 O
Galena	–	PbS		Epsom salt	– MgSO4 .7H 2 O
Copper pyrites			– CuFeS₂	Barytes	– BaSO₄
Cinnabar			– HgS	Zinc blende	– ZnS

Extraction of sulphur (Frasch process) : Sulphur is generally extracted from underground deposits by drilling three concentric pipes upto the beds of sulphur (700 – 1200 feet deep).
Allotropy in sulphur : Sulphur exists in four allotropic forms,

Rhombic or octahedral or a -sulphur : It is a bright yellow solid, soluble in All other varieties of sulphur gradually change into this form on standing.
Monoclinic sulphur or prismatic or b -sulphur: It is prepared by melting the sulphur and then cooling it till a crust is formed. On removing the crust, needle shaped crystals of monoclinic sulphur separate It is

CS2

and stable at room

dull yellow in colour, soluble in

CS2

and stable only above 369K. Below

this temperature it changes into rhombic form.

Thus, at 369K both these varities co-exist. This temperature is called transition temperature and the two sulphurs are called enantiotropic substances. It also exist as molecules similar to that of rhombic sulphur but the symmetry of the crystals is different.

Plastic or amorphous or g -sulphur : It is a super cooled liquid insoluble in It consists of long zig-zag chains of S-atoms.

CS₂ , soft and

Colloidal or d -sulphur : It is prepared by passing water or by treating sodium thiosulphate with HCl.

H 2 S

through a solution of an oxidizing agent or

Properties of sulphur : It burns in air with, a blue flame forming

SO₂ , gives sulphur hexafluoride with

F₂ and sulphur mono chloride with Cl₂ , sulphides with metals like Na, Ca, Zn, Hg, Fe, Cu etc., reduces HNO₃ to

NO2

and

H 2 SO4

to SO2 . With NaOH solution on heating, S8 + 12NaOH ¾¾® 4 Na2 S + 2Na2 S2 O3 + 6H 2 O . It

gives sodium sulphide and sodium thiosulphate, with excess of sulphur,

2Na2 S + S8 ¾¾® 2Na2 S5 .

Uses of sulphur : It is used in the manufacture of matches, gun powder (mixture of charcoal, sulphur

and potassium nitrate), explosives and fire works

SO2 , H 2 SO4 ,

CS2

and dyes, sulpha drugs and ointment for

curing skin diseases and in the vulcanization of rubber.

Compounds of Sulphur

Hydrogen Sulphide : It is prepared in the laboratory by the action of

H 2 SO4

on ferrous sulphide in

kipp’s apparatus,

FeS + H 2 SO4 ® FeSO4 + H 2 S . It is colourless gas having foul smell resembling that of rotten

eggs. It reacts with many cations (of group II and IV) to give coloured sulphides,

Cu⁺² + S ^-2 ® CuS ; Cd ⁺² + S ^-2 ® CdS ; Ni ⁺² + S ^-2 ® NiS ; Co ⁺² + S ^-2 ® CoS

Black

(Yellow)

(Black)

The solubility of sulphides can be controlled by the use in qualitative analysis of cation radicals.

H ⁺ ions concentration and therefore,

H ₂ S finds extensive

Oxides of sulphur : Sulphur forms several oxides of which sulphur dioxide (SO2 ) and sulphur trioxide

(SO₃ ) are most important.

Sulphur dioxide (SO₂) : It is prepared by burning sulphur or iron pyrites in

S8 + 8O2

® 8SO2 ;

4 FeS2

+ 11O2

® 2Fe

2 O3

+ 8SO2

In laboratory, it is prepared by heating copper turnings with conc.

Cu + 2H 2 SO4 ® CuSO4 + SO2 + 2H 2 O

H 2 SO4

It is a colourless gas with irritating and suffocating smell.

SO2

molecule has a bent structure with a O – S – O bond angle of 119^o. Sulphur is

sp ² hybridized.

Sulphur trioxide (SO₃) : It is formed by the oxidation of SO2 .

2SO2

+ O2

¾¾700¾K, 2¾at¾m. ® 2SO3

V2O5

In the gaseous phase, it exists as planar triangular molecular species involving hybridization of the S-atom. It has three S–O s bonds and three S–O p bonds. The O–S–O bond angle is of 120^o.

Oxyacids of sulphur : Sulphur forms many Some of these are,

*Formula*	Name	*Important properties*	*Structural formula*
H ₂ SO₃ (+4)	Sulphurous acid	Free acid does not exist	. . O = S – OH \| OH O \|\| O = S – OH \| OH S \|\| O = S – OH \| OH O O \|\| \|\| HO – S– S– OH O O \|\| \|\| O = S — S = O \| \| OH OH O O \|\| \|\| O = S – O– S = OH \| \| OH OH O \|\| HO = S — OOH \|\| O O O \|\| \|\| O = S – O – O – S = O \| \| OH OH
		diprotic, strong reducing
		agent
H ₂ SO₄ (+6)	Sulphuric acid	Stable diprotic, dehydrating
(Oil of vitriol)		agent
H ₂S₂O₃ (–2 and +6)	Thiosulphuric acid	Free acid does not exist but
		its salts e.g. Na2 S2 O3 All
		quite stable reducing agent
H2S2O4 (+3)	Dithionous acid
H 2 S2 O6 (+5)	Dithionic acid	Free acid is moderately
		stable but its salts are quite
		stable.
H 2 S2 O7 (+6)	Disulphuric acid	Strong oxidising agent
(Oleum)	(Pyrosulphuric acid)
H ₂ SO₅ (+6)	Peroxomonosulphuric acid	Stable crystalline solid,
(Caro’s acid)	(Its salts known as persulphates)	powerfull oxidising agent
H 2 S2 O8 (+6)	Peroxodisulphuric acid	Strong oxidising agent.
(Marshals acid)	(its salts are known as disulphates)

Sulphuric acid (H₂SO₄) : H₂SO₄ is a very stable oxyacid of It is often called king of chemicals, since it is one of the most useful chemicals in industry.

Manufacture of sulphuric acid : H₂SO₄ can be manufactured by following process,

Lead chamber process : In this process, SO₂ is oxidized to SO₃ by the oxides of nitrogen and the SO₃ thus formed is dissolved in steam to form H₂SO₄.

SO₂ + NO₂ ® SO₃ + NO ; 2NO + O₂ ® 2NO₂; SO₃ + H₂O ® H₂SO₄

Contact process : In the contact process, SO₂ obtained by burning of S or iron pyrities is catalytically oxidized to SO₃ in presence of finely divided Pt or V₂O₅ as

V₂O₅ or Pt, 673-732 K

S + O₂ ® SO₂ or 4FeS₂ + 11O₂ ® 2Fe₂O₃ + 8SO₂ ; 2SO2 + O2 2SO3 .

V₂O₅ is, however, preferred since is much cheaper than Pt and is also not poisoned by arsenic impurities. The favorable conditions for maximum yield of SO₃ are,

High concentration of SO₂ and O (ii) Low temperature of 673 to 723 K, (iii) High pressure about 2 atmospheres.

SO₃ thus obtained is absorbed in 98% H₂SO₄ to form oleum which on dilution with water gives H₂SO₄ of desired concentration.

SO3 + H 2 SO4

® H₂S₂O₇ ;

oleum

H 2 S2 O7 + H 2 O ® 2H 2 SO4

Contact process is preferred over lead chamber process (gives 98% pure greater purity (100%).

H 2 SO4 ) since it gives H₂SO₄ of

Flow sheet diagram of it’s preparation is as follows

H₂SO₄	Dilution	Oleum		Absorption	SO₃	Catalyst		Pre heater
				Tower		Chamber

Structure : H₂SO₄ is a covalent compound and has tetrahedral (S is sp³– hybridized) structure.

Properties : H₂SO₄ has high b.p. (611K) and is also highly viscous due to H-bonding. It has strong affinity for H₂O and a large amount of heat is evolved when it is mixed with water.

H₂SO₄ is a strong dibasic It neutralizes alkalies, liberates CO₂ from carbonates and bicarbonates.
It reacts with more electropositive (than hydrogen) metals to evolve H₂ and produces SO₂ on heating with less electropositive metals than hydrogen .eg.,

H 2 SO4 + 2KOH ® K2 SO4 + 2H 2 O ; Cu + 2H 2 SO4 ® CuSO4 + SO2 + 2H 2 O

It is a strong oxidizing agent and oxidises as follows,

H 2 SO4 ® H 2 O + SO2 + O ;

S + 2H 2 SO4 ® 3SO2 + 2H 2 O ;

C + 2H 2 SO4 ® 2SO + CO + 2H 2 O

P4 + 10H 2 SO4 ® 4 H 2 PO4 + 10SO2 + 4 H 2 O

2HBr + H 2 SO4 ® Br2 + 2H 2 O + SO2 ; 2HI + H 2 SO4 ® 2H 2 O + I 2 + 2SO2

It reacts with number of It liberates HCl from chlorides, H ₂ S from sulphides, HNO₃ from nitrates.
It acts as a strong dehydrating agent, as it dehydrates, sugar to sugar charcoal (carbon), formic acid to CO, oxalic acid to CO+ CO₂ and ethyl alcohol to
It is also a good sulphonating agent and used for sulphonation of aromatic compounds. ,

BaCl 2 + H 2 SO4 ® BaSO4 + 2HCl ;

(white ppt)

Pb(NO3 )2 + H 2 SO4 ® PbSO4 + 2HNO3

(white ppt.)

C12 H 22 O11 ¾¾^Con¾^c. ^H¾2SO¾4 ® 12C + 11H 2 O ; HCOOH ¾¾Con¾c.H¾2 SO¾4 ®

CO + H₂O

Sugar

Carbon

Uses : H₂SO₄ is used (i) in the preparation of fertilizers like (NH₄)₂ SO4 and super phosphate of lime, (ii) in

lead storage batteries (iii) in preparation of dyes, paints and explosives (iv) in textile and paper industry (v) for training of tanning (vi) as a dehydrating agent.

(5) Sodium thiosulphate

Na2 S2 O3 .5H 2 O : It is manufactured by saturating a solution of sodium carbonate

with SO₂ which gives a solution of sodium sulphite, Na2 CO3 + SO2 + H 2 O ® Na2 SO3 + CO2 + H 2 O

The resulting solution is boiled with powdered sulphur as,

Na₂SO₃ + S ¾¾³⁷³¾^K ®

Na 2 S2 O3

The solution is then cooled to get crystals of sodium thiosulphate.

Physical properties : (1) Sodium thiosulphate is a colourless crystalline solid. In the hydrated form, it is called hypo. (2) It melts at 320 K and loses its water molecules of crystallization on heating to 490K.

Chemical properties

Action with halogens : It reacts with halogens as,
- Chlorine water oxidizes sodium thiosulphate to sodium sulphate and sulphur is precipitated,

Na ₂S₂O₃ + Cl₂ + H₂O ® 2HCI + Na₂SO₄ +S

This property enables it to act as an antichlor in bleaching i.e. it destroys the unreacted chlorine in the process of bleaching.

Bromine water also oxidizes sodium thiosulphate to sodium sulphate and sulphur,

Na ₂S₂O₃ + Br₂ + H₂O ® Na₂SO₄ + 2HBr + S

With iodine it forms a soluble compound called sodium tetrathionate,

2Na2 S2 O3 + I 2 ®

Na2 S4 O6

Sod. tetrathionate

2Nal

Therefore, hypo is commonly used to remove iodine stains from the clothes.

Action of heat : Upon heating, sodium thiosulphate decomposes to form sodium sulphate and sodium

pentasulphide,

4 Na 2 S2 O3 ¾¾^He¾^at ® 3Na 2 SO4 +

Na 2 S5

Sodium pentasulphide

Action with acids : Sodium thiosulphate reacts with dilute hydrochloric acid or Sulphuric acid forming sulphur dioxide and The solution turns milky yellow due to sulphur.

Na₂S₂O₃ + 2HCI ® 2NaCl + SO₂ + H₂O + S

Action with silver halides : Sodium thiosulphate forms soluble complex when treated with silver

chloride or silver bromide,

2Na2 S2 O3 + 2AgBr ® Na3 Ag(S2 O3 )2 + NaBr .

Sodium dithiosulphate argentate (I) compex

This property of hypo is made use in photography.

Uses of sodium thiosulphate

It is largely used in photography as a fixing
It is used as a preservative for fruit products such as jams and
It is used as an antichlor in
It is used as a volumetric agent for the estimation of

It is used in

Fluorine is the first member of group 17 or VIIA of the periodic table. It consists of five elements Fluorine (F), Chlorine (Cl), bromine (Br), iodine (I) and astatine (At). These are known as halogen because their salts are found in sea water. Halogen is a greek word meaning a sea salt.

Electronic configuration

Elements Electronic configuration (

ns ² np⁵

)

9 F 1s ² ,2s ² 2p⁵ or [He] 2s ² 2p⁵

17 Cl 1s ² ,2s ² 2p⁶ ,3s ² 3p⁵ or [Ne] 3s ² 3p⁵

35 Br 1s ² ,2s ² 2p⁶ ,3s ² 3p⁶ 3d¹⁰ ,4s ² 4 p⁵ or [Ar] 3d¹⁰ 4s ² 4 p⁵

53 I 1s ² ,2s ² 2p⁶ ,3s ² 3p⁶ 3d¹⁰ ,4s ² 4 p⁶ 4d¹⁰ ,5s ² 5p⁵ or [Kr] 4d¹⁰ 5s ² 5p⁵

85 At 1s ² ,2s ² 2p⁶ ,3s ² 3p⁶ 3d¹⁰ ,4s ² 4 p⁶ 4d¹⁰ 4 f ¹⁴ ,5s ² 5p⁶ 5d¹⁰ ,6s ² 6p⁵ or [Xe] 4 f ¹⁴ 5d¹⁰ 6s ² 6p⁵

Physical properties

Physical state : Halogens exist as diatomic covalent

F2 – gas, Cl 2 – gas, Br2 – corrosive liquid, I 2 – volatile solid.

Melting and boiling points : They increase with increase in atomic
Ionization energy : The E. decreases on moving down the gorup.
Electron affinity : F < Cl > Br > I or Cl > F > Br >
Electronegativity : F > Cl > Br >

(6) Bond energy

Element	F – F	Cl – Cl	Br – Br	I – I
Bond length (Å)	1.42	1.99	2.28	2.67
Bond dissociation energy (kcal / mole)	38	57	45.5	35.6

Colour : F – Light yellow , Cl – Greenish yellow, Br – Reddish brown, I – Deep

Oxidation state : All exhibit –1 Oxidation state Except fluorine other element also show +3 ,+5, +7 oxidation
Oxidising power : F2 > Cl 2 > Br2 > I 2 .

Solubility : Halogen being non polar in nature do not dissolve in water

2F2 + 2H 2O ® 4HF + O2 ,

3F2 + 3H 2O ® 6HF + O3 (fluorine highly soluble) Cl 2 and

Chemical properties

Br₂ are fairly soluble. I ₂ is a least soluble in water.

Reactivity : The halogen are most reactive elements due to their low bond dissociation energy, high electron affinity and high enthalpy of hydration of halide F > Cl > Br > I
Reaction with H₂O : Halogens readily decomposes water. This tendency decreases on moving down the Fluorine decomposes water very energetically to give oxygen and ozone,

2H ₂O + 2F₂ ® 4 HF +

O2 ;

Oxygen

3H ₂O + 3F₂ ® 6HF +

Ozone

Fluorine gives fumes in moist air. This is due to the formation of HF, which is a liquid and can absorb moisture to form liquid droplets and therefore, gives fumes with moist air. Chlorine and bromine react less

vigorously, Cl2 + H 2 O ® HCl +

HClO ;

Hypochlorous acid

Br₂ + H ₂O ® HBr +

HBrO

Hypobromous acid

In the presence of sunlight, HClO (hypochlorous acid) HBrO (hypobromous acid) liberate oxygen.

2HClO ® 2HCl + O2 ; 2HBrO ® 2HBr + O2

Iodine is only slightly soluble in water. However, it dissolves in 10% aqueous solution of Kl due to the

formation of

I 2 + KI ⇌

KI 2 or

I 2 + I ^– ⇌

–

Complex ion

Reaction with hydrogen : Form covalent

2 2

H + F ¾¾^–²⁰¾⁰o¾^C ® 2HF (very violent); H + Cl ¾¾^Sunli¾^g¾^h¾^t ® 2HCl

2 2

H 2 + Br2

¾¾^H^e¾^a¾^t ® 2HBr ;

catalyst

H 2 + I 2

Heat 2HI

(poor yield)

Acidic strength in aqueous solution is in the order, HI > HBr > HCl < HF.

Reducing character of hydrides follow the order, HI > HBr > HCl > HF.

Boiling point HF > HI > HBr > HCl. Thermal stability, H – F > H – Cl > H – Br > H – I. HCl is also called Muriatic acid.

Hydrides : All the halogens combine directly with hydrogen to form halogen acids but their reactivity progressively decreases from fluorine to iodine, H₂ + X₂ ® 2HX (X = F, Cl, Br or I).
Boiling points or volatility : In other words volatility decreases in the order : HCl > HBr > HI > HF as the boiling points increase in the order : HCl (189K) < HBr (206K) < HI (238K) < HF (292.5K).
Thermal stability : Thermal stability of the hydrides decrease from HF to HI e., HF > HCl > HBr > HI.
Acidic strength : The acidic strength of halogen acids decreases from HI to HF e, HI > HBr > HCl > HF.

Reducing properties : Since the stability of hydrides decreases from HF to Hl, their reducing properties increase in the order HF < HCl < HBr < H.
Dipole moments : The dipole moments of hydrogen halides decrease in the order : HF > HCI > HBr > HX

as the electro negativity of the halogen atom decreases form F to I.

HX HF HCl HBr Hl

Dipole moment (D) 1.74 1.07 0.78 0.38

Oxides : Halogens (except F₂) do not combine readily with However, a number of compounds of halogens with oxygen have been prepared by indirect methods. Only two compounds of fluorine with oxygen,

i.e. oxygen difluorine (OF₂) and oxygen fluoride (O₂F₂) are known. Chlorine forms largest number of oxides i.e. Cl₂O, ClO₂, Cl₂O₆ and Cl₂O₇ while iodine forms the least, i.e. I₂O₅. Bromine, however, forms three oxides (Br₂O, BrO₂C BrO₃). In all these compounds, bonds are largely covalent. All the oxides of halogens are powerful oxidizing agents. These compounds are very reactive and are unstable towards heat. The stability of oxides is greatest for iodine while bromine oxides are the least stable. For a particular halogen, higher oxides are more stable than the lower ones.

Iodine-oxygen bond is stable due to greater polarity of the bond (due to larger electro negativity difference between I and O) while in chlorine-oxygen bond, the stability is gained through multiple bond formation involving the d-orbital of chlorine atom. Bromine lacks both these characteristics and hence forms least stable oxides.

Oxides of chlorine, bromine and iodine are acidic and the acidic character increases as the percentage of oxygen increases in them.

Iodine also forms l₂O₄ and l₄O₉ compounds which are believed not to be true oxides but are basic iodyliodate, IO (IO₃) and normal iodine triodate, I (IO₃)₃ having tripositive iodine as the cation.

OF₂ is V-shaped having bond angle 103^o, Cl₂O is also V-shaped with bond angle 111^o while ClO₂ is angular with-bond angle 118^o. It is paramagnetic due to odd number of electrons having three-electron bond. It is regarded as a mixed anhydride of chloric and chlorous acids. 2ClO₂ + H₂O ® HClO₂ + HClO₃

Oxoacids of halogens : Fluorine does not form any oxoacid since it is the strongest oxidizing agent. Chlorine, bromine and iodine mainly form four series of oxoacids namely hypohalous acid (HXO), halous acid (HXO₂) halic acid (HXO₃) and perhalic acid (HXO₄) as given below :

*Oxidation state*	*Chlorine*	*Bromine*	*Iodine*	*Thermal* *stability* *and acid strength*	*Oxidising power*
+1	HClO	HBrO	HIO
+3	HClO₂	–	–
+5	HClO₃	HBrO₃	HIO₃
+7	HClO₄	HBrO₄	HIO₄
	Acidity decreases ®

Hybridized ion : In all these oxoacids, the halogen atom is sp³ -hybridized.

Acidic character : All these acids are monobasic containing an—OH The acidic character of the oxoacids increases with increase in oxidation number, i.e., HClO < HClO₂ < HClO₃ < HClO₄ and the strength of

the conjugate bases of these acids follows the order, ClO ^– > ClO ^– ® ClO ^– > ClO^– .

2 3 4

Oxidising power and thermal stability : The oxidizing power of these acids decreases as the oxidation number increases, i.e., HClO < HClO₂ < HClO₃ < HClO₄. Stability of oxoacids of chlorine in the increasing order is , HClO < HClO₂ < HClO₃ < HClO₄ and the increasing stability order of anions of oxoacids of

chlorine is, ClO ^– < ClO ^– < ClO ^– < ClO ^– .

2 3 4

As the number of oxygen atoms in an ion increases there will be a greater dispersal of negative charge and thus greater will be the stability of ion formed. For different halogen having the name oxidation number, the thermal stability decreases with increase in atomic number i.e., it is in the order HClO > HBrO > HIO and ClO^– >

BrO^– > IO^– However, in

HXO3

is most stable. The stability order being HClO₃ < HBrO₃ < HIO₃.

Perhalates are strong oxidizing agents, the oxidizing power is in the order, BrO^– > IO^– > ClO^– .

4 4 4

Thus BrO₄ is the strongest oxidizing agent (though its reaction is quite slow) and ClO ^–

is the weakest.

The acidity of oxoacids of different halogens having the same oxidation number decreases with increase in

the atomic size of the halogen i.e. HClO₄

HBrO4

HIO4 .

Reaction with alkalies : 2F2 + 2NaOH ® 2NaF + OF2 + H 2 O

(cold dilute)

; 2F + 4 NaOH ® 4 NaF + O2 + 2H 2 O

(hot conc.)

Halogen other than fluorine (Cl 2 ,Br2 , I 2 ) react with NaOH as follows,

X 2 (g) +

2OH ^–

(cold dilute)

¾¾¹⁵o¾^C ® X ^– + OX ^– + H O ;

X (g ) + 6OH ^– ¾¾⁷⁰o¾^C ® 5X ^– +

(hot conc)

XO ^–

+ 3H 2O

(hypohalite ion) (halate ion)

Bleaching action of halogen : Cl ₂ acts as bleaching agent, its bleaching action is permanent Cl ₂ water can also act as ink

(9) Reaction with other halides

2KBr(aq.) + Cl₂ (g) ® 2KCl(aq.) + Br(aq.) ; 2KI (aq.) + Cl ₂ (g ) ® 2KCl(aq.) + I ₂ (aq.)

Inter halogen compounds : The compounds of one halogen with the other are called interhalogens or interhalogen compounds. The main reason for their formation is the large electronegativity and the size differences between the different Taking A as the less electronegative and B as the more electronegative halogen, they are divided into the following four types the less electronegative halogen (A) is always written first.

AB₃

AB₅

AB₇

ClF

BrF, BrCl, ICl

IBr, IF

ClF₃ , BrF₃

IF3 , ICl3

BrF5 IF5

IF7

These interhalogen compounds are unstable and more reactive

(i) General properties

Largest halogen always serves the central
The highest interhalogen compound e. lF₇ is obtained with iodine, the largest halogen attached to the smallest one

The bonds in interhalogen compounds are essentially
Thermal stability decreases as the size difference decreases and increases as the polarity of the bond Thus ClF is thermally more stable as compared to IBr.

They ionize in solution or in the liquid state,

2ICl

⇌ I ⁺ + ICl ^– ;

2ICl ⇌ ICl ⁺ + ICl ^–

Hydrolysis of interhalogen compounds always produces a halide ion derived from smaller halogen and

oxyhalide derived from larger halogen,

ICl + H 2 O ® Cl ^– + OI ^– + 2H ⁺ ;

BrF5 3H 2 O ® 5F ^– + BrO^– + 6H ⁺

They are strong oxidizing
Largest number of interhalogens are formed by fluorine due to its smaller size and higher electronegativity or oxidizing
Structure : Interhalogen compounds are,
AB type e. ICl, IBr, IF etc, are linear
AB₃ type e. IF₃, ClF₃, BrF₃ have distorted trigonal bipyramidal (dsp³-hybridization) structures of T-shape due to two lone pairs in equatorial positions ICl₃ is dimeric, I₂Cl₆ and has a planar structure.
AB₅ types e. BrF₅, IF₅ have distorted octahedral (d²sp³-hybridization) shapes or square pyramidal due to a lone pair one of the axial positions.
AB₇ type e. IF₇, have pentagonal bipyramidal (d³sp³-hybridization) structures.

(11) Polyhalides :

KI + I ₂ ® KI ⇌ K ⁺ + I ^– ;

Cl ^– , Br ^– , I ^– , ICl ^– , IBr ^– , ICl ^– , BrF ^– , I ^– , IF ^–

and I ^–

(12) Pseudohalogen and pseudohalides

3 3 3 2 2 4

4 5 6 7

*Pseudohalogen*	*Pseudohalide*
Cyanogen – (CN)₂	Cyanide – CN ^–
Oxocyanogen – (OCN)₂	Cyanate – OCN ^–
Thiocyanogen – (SCN)₂	Thiocyanate – SCN ^–
Selenocyanogen – (SeCN)₂	Selenocyanate – SeCN ^–

Freons : Freon –11 is

CCl3 F ; Freon –12 is CCl 2F2

and it is marketed under the popular brand names

such as ‘Freon’ and ‘Genetron’ ; Freon –113 is

CClF2 . CF3 . These cause ozone depletion.

Preparation of halogens

CCl2 F. CClF2 ; Freon –114 is

CClF2 . CClF2 ; Freon –115 is

Preparation of fluorine : F₂ is prepared by electrolysis of a solution of KHF₂ (1 Part) in HF (5 part) in a vessel (Modern method) made of Ni – Cu alloy or Ni –Cu– Fe alloy called the monel metal using carbon

During the electrolysis following reactions occur,

KHF2 ⇌

KF + HF ; KF ⇌ K ⁺ + F ^– .

At cathode : K⁺ + e^– ® K;

2K + 2HF ® 2KF + H ₂ ; At anode :

F ^– ® F + e – ;

F + F ® F2

Preparation of chlorine : On the industrial scale, Cl₂ is prepared by the electrolysis of concentrated aqueous solution of NaCl. In this process, NaOH and H₂ are by

2 NaCl (aq) +2H₂O ¾¾^El^e^c¾^tr^ic¾^it^y ® 2NaOH(aq) +Cl₂ +H₂

In the laboratory, Cl₂ can be prepared by adding conc HCl on KMnO₄ or MnO₂.

2KMnO₄ + 16HCl ® 2KCl + 2MnCl₂ + 8H₂O + 5Cl₂ ; MnO₂ + 4HCl ® MnCl₂ + 2H₂O + Cl₂

Preparation of Bromine : In laboratory it is prepared by heating NaBr with MnO₂ and Conc H₂SO₄. 2NaBr + MnO₂ + 3H₂SO₄ ® 2NaHSO₄ + MnSO₄ + 2H₂O + Br₂

Preparation of Iodine (Lab method) : By heating a mixture of

2KI + H ₂ SO₄ ¾¾^Mⁿ^O¾2 ® K ₂ SO₄ + SO₂ + I ₂ + 2H ₂O . 2I ^– + Cl₂ ® I₂ + 2Cl.

I₂ is commercially prepared from sea weeds.

Uses of Halogens

MnO2 , H 2 SO4

and an iodide

Uses of Fluorine : (i) It is used as an oxidising agent and fluorinating agent. (ii) Fluorine and its compounds such as NF₃, OF₂ are used as rocket fuels. (iii) It is used in the manufacture of a plastic known as Teflon (CF₂-CF₂)n which is resistant to the action of all acids, alkalies and even boiling aqua regia. (iv) It is used in the manufacture of fluorocarbons like freon which is used as an excellent refrigerant and in air (v) It is used for the preparation of uranium hexafluoride, which is used for the separation of isotopes of U(235) and U(238).
Uses of Chlorine : (i) Chlorine is used in sterilization of drinking (ii) Large quantities of chlorine are used industrially for the bleaching of cotton, paper, wood, textiles, etc. (iii) It is used in making insecticides like D.D.T., germicides, dyes, drugs, etc. (iv) It is used for preparing vinyl chloride which is a starting material for making th plastic PVC. (v) It is used in the manufacture of chlorinated organic solvents like CHCl₃, CCl₄, which are used for dry cleaning and degreasing machinery. (vi) It is used in the preparation of HCl, bleaching powder, chlorates, perchlorates, sodium hypochlorite which are important industrial compounds.
Uses of Bromine : (i) Bromine is used in the preparation of ethylene bromide, which is mixed with tetraethyl lead (TEL) and added to the petrol as an anti-knocking (ii) In the manufacture of AgBr used in photography. (iii) In the manufacture of dyes, drugs, etc. (iv) It is used in the manufacture of benzyl bromide which is an effective teargas. (v) It is used as a laboratory regent.

Uses of Iodine : (i) Iodine is used as a laboratory (ii) It is used in making medicines and dyes. Tincture of iodine is an antiseptic. (iii) AgI is used in photographic emulsions. (iv) It is used in the preparation of iodized salt. Iodized salt is used to prevent the occurrence of common goiter.

Helium is the first member of group 18 or zero of the periodic table. It consists of six elements helium (He), Neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn). Zero group occupies the intermediate position between the elements of VIIA (17th) and IA (1st) groups. These are collectively called as inactive gases or inert gases. However, these are now called noble gases as some compounds of these gases have been obtained under certain specific conditions.

Electronic configuration

Elements Electronic configuration (

ns ² np⁶

)

2 He 1s ²

10 Ne 1s ² ,2s ² 2p⁶

18 Ar 1s ² ,2s ² 2p⁶ ,3s ² 3p⁶

36 Kr 1s ² ,2s ² 2p⁶ ,3s ² 3p⁶ 3d¹⁰ ,4s ² 4 p⁶

Atomic radii : The atomic radii of noble gases increases on moving down the group and their atomic radii correspond to the van der Waal’s
Boiling points : The pt. and b.pt. increases from He to Rn, because of increase in magnitude of Vander Waal’s forces.
Polarizabiltiy : The polarizability increases down the group, He < Ne < Ar < Kr < Xe

Ionisation energy and electron affinity : Noble gases have stable

ns ²np⁶

fully filled electronic

configuration, so these have no tendency to add or lose electron. Therefore, ionisation energy of noble gases is very high. On the other hand their electron affinity is zero.

Heat of vaporisation : They posses very low values of heat of vapourisation, because of presence of very weak Vander Waal’s forces of attraction between their monoatomic However the value of heat of vaporisation increases with atomic number down the group and this shows that there is an increasing polarizability of the larger electronic clouds of the elements with higher atomic number.
Solubility in water : They are slightly soluble in water. Their solubility generally increases with the increase in atomic number down the
Adsorption by charcoal : All of them except helium are adsorbed by cocount charcoal at low The extent of adsorption increases down the group.
Characteristic spectra : All of them give characteristic spectra, by which they can be
Liquification of gases : It is difficult to liquify noble gases as their atoms are held by weak Vander Waal’s forces. Ease of liquification increases down the group from He to Rn. Helium has the lowest boiling point (4.18 K) of any known substance. The ease of liquification increases down the group due to increase in intermolecular

The elements helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn), constitute zero group of the periodic table. These are gases at ordinary temperature and do not have chemical reactivity and therefore, these are called inert gases.

Occurrence : Due to the inert nature of noble gases, they always occur in the free state. Except radon, all these gases are present in atmosphere in the atomic

Element	He	Ne	Ar	Kr	Xe
Abundance (Volume %)	5.2 ´ 10^–⁴	1.8 ´ 10^–³	9.3 ´ 10^–¹	1.4 ´ 10^–³	8.7 ´ 10^–⁶

He is also present in natural gas to the extent of 2 to 7%.

(11) Isolation

Helium : It is commercially obtained from natural gas. The natural gas contains hydrocarbons (methane ), CO₂, H₂S and He as the main constituents.

The natural gas is compressed to about 100 atm and cooled to 73K. He remains unliquefied while other gases get liquefied. About 99% pure He is prepared by this method.

Argon, Neon, Krypton and Xenon : These gases are prepared by the fractionation distillation of liquid Fractional distillation of air gives O₂, N₂ and mixture of noble gases. The individual gases may be obtained by adsorption of air on coconut charcoal. The charcoal adsorbs different gases at different temperatures and can be collected.
Radon : It can be obtained by radio active disintegration of radium (226), ₈₈Ra²²⁶ ® ₈₆Rn²²²+ ₂a ⁴ .

(12) Compounds of Xenon

In 1962, N. Bartlett noticed that PtF₆ is a powerful oxidizing agent which combines with molecular oxygen to

form ionic compound, dioxygenyl hexafluoro platinate (v) O⁺ [PtF6 ]^– , O2(g) + PtF6(g) ® O ⁺ [PtF6 ]^– ,This indicates

2 2

that PtF₆ has oxidized O₂ to O ⁺ . Now, oxygen and xenon have some similarities,

The first ionization energy of Xe gas (1170kJ mol ^-1 ) is fairly close to that of oxygen (1166kJ mol ^-1 ) .
The molecular diameter of oxygen and atomic radius of Xe are similar (4Å)

On this assumption, Bartlett reacted Xenon and

PtF6

in gas phase and a orange yellow solid of the

composition Xe PtF₆ was obtained,

Xe(g) + PtF6(g) ® Xe ⁺ [PtF6

–

]

(s)

Some important stable compounds of Xe are,

Orange yellow

+2	+4		+6
XeF2	XeF4 ,	XeOF2	XeF6 ,	XeOF4 ,	XeO3

Fluorides : Xenon forms three compounds with fluorine. These are : Xenon difluoride (XeF2), Xenon tetrafluoride (XeF₄) and Xenon hexafluoride (XeF₆).

Xenon difluoride (XeF₂) is formed when a mixture of Xenon and fluorine in the ratio 1 : 3 by volume is

passed through a nickel tube at 673 K,

Xe + Fe ¾¾^Ni, ⁶¾⁷³¾^K ® XeF2

Structure : XeF₂ has trigonal bipyramid geometry due to

sp ³

d-hybridization of Xe. Three equatorial

positions are occupied by lone pairs of electrons giving a linear shape to the molecule.

Properties : XeF₂ is a colourless crystalline solid, reacts with H ₂

to give Xe and HF. It is hydrolysed

completely by water ,

2XeF2 + 2H 2 O ® 2Xe + O2 + 4 HF .

It also forms addition compounds with reactive pentafluorides like SbF₅, TaF₅ etc.

XeF2 + 2SbF5 ® XeF2 . 2SbF5

It is a mild fluorinating agent and hence reacts with benzene to give fluorobenzene.

Xenon tetrafluoride (XeF₄) is prepared by heating a mixture of xenon and fluorine in the ratio 1 : 5 in a

nickel vessel at 673 K and then suddenly cooling it in acetone.

XeF4

is also formed when an electric discharge is

passed through a mixture of xenon and excess of fluorine,

Xe + 2F2 ¾¾^Ni, ⁶¾⁷³¾^K ® XeF4

Structure :

XeF₄ has square planar shape due to

sp ³d ²

hybridizationof Xe giving

octahedral geometry with two trans positions occupied by lone pairs of electrons.

Properties :

XeF4

is a colourless, crystalline solid, soluble in anhydrous HF, reacts with

H ₂ to form Xe and HF and reacts with water to give highly explosive solid, 6 XeF4 + 12H 2 O ® 4 Xe + 2XeO3 + 24 HF + 3O2

XeO3 . (complete hydrolysis),

Partial hydrolysis yields XeOF₂,

XeF4 + H2O ¾¾¹⁹³¾^K ® XeOF2 + 2HF

It also forms addition compounds with SbF₅, It also acts as a strong fluorinating agent.

XeF4 + SbF5 ® [XeF3 ]⁺[SbF6 ]^– .

Xenon hexafluoride (XeF₆) is prepared by heating a mixture of xenon and fluorine in the ratio 1 : 20 at

473—523 K under a pressure of 50 atmospheres.

Xe + 3F2 ¾¾473¾- 52¾3K¾, 50¾at¾m. ® XeF6

Structure : XeF₆ has pentagonal bipyramid geometry due to sp³d³ hybridization. One trans position is occupied by a lone pair giving a distorted octahedral shape.

Properties : It is colourless, crystalline solid, highly soluble in anhydrous HF giving

solution which is a good conductor of electricity,

HF + XeF₆ ® XeF ⁺ + HF ^– .

It is the most powerful fluorinating agent and reacts with H₂ to give Xe and HF. Partial hydrolysis of XeF₆ yields XeOF₄ an complete hydrolysis yields xenon trioxide, XeO₃.

XeF6 + H2O ® XeOF4 + 2HF; XeF6 + 3H 2 O ® XeO3 + 6HF

It forms addition compounds with alkali metal fluorides (except LiF) of the formula XeF₆. MF where M represents the alkali metal.

Oxides : Xenon forms two oxides such as xenon trioxide (XeO₃) and xenon tetraoxide (XeO₄).

Xenon trioxide (XeO₃) is prepared by complete hydrolysis of XeF₄ and XeF₆

6 XeF4 + 12H 2 ® 2XeO3 + 4 Xe + 3O2 + 24 HF ; XeF6 + 3H2O ® XeO3 + 6HF

Structure : XeO₃ has tetrahedral geometry due to

sp³

hybridization of Xe. One of the hybrid orbitals

contains a lone pair of electrons giving a trigonal pyramidal shape. The molecule has three Xe = O double bonds

containing

pp – dp

overlapping.

Properties : It is a colourless solid, highly explosive and powerful oxidizing agent.

Xenon tetraoxide (XeO₄) is prepared by the action of H₂SO₄ on sodium or barium xenate

(Na 4 XeO6 ; Ba2 XeO6 ) at room temperature,

Na4 XeO6 + 2H 2 SO4 ® XeO4 + 2Na2 SO4 2H 2 O ;

XeO₄ is purified by vacuum sublimation at 195 K.

Structure : XeO₄ has tetrahedral structure due to

Ba2 XeO6 + 2H 2 SO4 ® XeO4 + 2BaSO4 + 2H 2 O

sp³ hybridization of Xe. There are four Xe–O double

bonds containing

pp – dp

overlapping.

Properties : It is quite unstable gas and decomposes to xenon and oxygen,

XeO₄ ® Xe + 2O₂ .

Oxyfluorides : Xenon forms three types of oxy fluorides such as xenon oxydifluoride (XeOF₂), xenon

oxytetrafluride

XeOF4

and xenon dioxydifluoride (XeO₂F₂).

Xenon oxydifluoride (XeOF₂) is formed by partial hydrolysis of XeF₄ at 193 K,

XeF4 + H 2 O ¾¾¹⁹³¾^K ® XeOF2 + 2HF .

Structure : XeOF₂ has trigonal bipyramid geometry due to sp³ d-hybridization of Xe. Two equatorial

positions are occupied by lone pairs of electrons giving a T-shape to the molecule. There is one Xe–O double bond

containing

pp – dp

overlapping.

Xenon oxytetrafluoride (XeOF₄) is prepared by partial hydrolysis of XeF₆; XeF6+ H 2 O ® XeOF4 + 2HF .

It can also be prepared by the reaction of SiO₂ with XeF₆,

Structure : XeOF₄ has octahedral geometry due to

2XeF6 + SiO2 ® 2XeOF4 + SiF4 .

sp³ d ² hybridization of Xe. One trans position is

occupied by a lone pair giving pyramid shape to the molecule. There is one Xe–O double bond containing

pp – dp overlapping.

Properties : It is a colourless volatile liquid which melts at 227 K. It reacts with water to give XeO₂F₂ and

XeO₃,

XeOF₄ + H₂O ® XeO₂ + 2HF ,

XeO2 F2 + H2O ® XeO3 + 2HF .

It is reduced by H₂ to Xe, XeOF4 + 3H2 ® Xe + H2O + 4 HF

Xenon dioxydifluoride (XeO₂F₂) is formed by partial hydrolysis of XeOF₄ or XeF₆

XeOF4 + H2O ® XeO2 F2 + 2HF ; XeF6 + 2H2O ® XeO2 F2 + 4 HF

It can also be prepared by mixing XeO₃ and XeOF₄ at low temperature (195K). The product is purified by

fractional distillation,

XeO3 + XeOF4 ¾¾¹⁹⁵¾^K ® 2XeO2 F2

Structure : XeO₂F₂ has trigonal bipyramid geometry due to sp³ d-hybridization of Xe. One equatorial

position is occupied by a lone pair of electrons giving a see-saw structure (shape) to the molecule. There are two

Xe–O double bonds containing

pp – dp

overlapping.

Properties : It is a colourless solid which melts at 303 K. It is easily hydrolysed to give XeO₃ XeO2 F2 + H2O ® XeO3 + 2HF

(13) Uses of noble gases

He is used for filling of balloons and air ships because of its non-inflammability and high power (which is 6% to that of hydrogen).
Oxygen-helium (1 : 4) mixture is used for treatment of asthma and for artificial respiration in deep sea diving because unlike nitrogen, helium is not soluble in blood even under high
Helium is also used for creating inert atmosphere in chemical
Liquid helium is used as a cryogenic fluid to produce and maintain extremely low temperatures for carrying out researches and as a coolant in atomic reactors and super conducting
It is also used in low temperature gas thermometry and as a shield gas for arc
Argon is used for creating inert atmosphere in chemical reactions, welding and metallurgical operations and for filling in incandescent and fluorescent It is also used in filling Geiger-Counter tubes and thermionic tubes.

Krypton and xenon are also used in gas filled A mixture of krypton and xenon is also used in some flash tubes for high speed photography.
Radon is used in radioactive research and therapeutics and in the non-surgical treatment of cancer and other malignant

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Electronic configuration

(10) Allotropy

Chemical properties

(1) Hydrides :

(3) Oxyacids :

Oxygen and its compounds

Physical properties of atomic and molecular oxygen

(5) Uses of dioxygen

(6) Compounds of Oxygen

Ozone or trioxygen

Structure of O3

Uses of ozone

Sulphur and its compounds

Compounds of Sulphur

(5) Sodium thiosulphate

Chemical properties

Uses of sodium thiosulphate

Electronic configuration

(6) Bond energy

Chemical properties

(9) Reaction with other halides

(i) General properties

(11) Polyhalides :

(12) Pseudohalogen and pseudohalides

Preparation of halogens

Uses of Halogens

Electronic configuration

(11) Isolation

(12) Compounds of Xenon

(13) Uses of noble gases

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Structure of O₃