Chapter 12 s & p Block Elements part 3 by TEACHING CARE Online coaching and tuition classes
- 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 ,
CaCN 2 ,
- 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.
|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 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
|Oxygen –||O2 and O3|
|Sulphur –||Rhombic , monoclinic, plastic sulphur|
|Selenium–||Red (non-metallic) grey (metallic)|
|Tellurium||–||Non-metallic and metallic (more|
|Polonium||–||a and b (both metallic)|
- Oxidation states : Oxygen ® – 2 and –1 oxidation These element shows +2 ,+4 and +6 oxidation state.
(1) Hydrides :
H 2 O, H 2 S, H 2 Se, H 2 Te
H 2 Po, H 2 O – colourless and odourless.
H 2 S, H 2 Se, H 2 Te
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
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 :
( Abundance: 99.76%)
Out of the three isotopes, 8 O18
( Abundance: 0.037%)
( 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 ¾¾He¾at ® 2KNO2 + O2 ; 2KClO3 ¾¾He¾at ® 2KCl + 3O2
- By heating dioxides, Peroxides and higher oxides : g.
2Ag 2 O ¾¾He¾at ® 4 Ag + O2 ;
¾¾He¾at ® Mn3 O4 + O2 ;
- Laboratory Method : In the laboratory, O2 is prepared by thermal decomposition of potassium
2KClO3 ¾¾420¾K ® 2KCl + 3O2
Note : ® In the absence of
catalyst, the decomposition takes place at 670-720 K. Therefore, MnO2
acts as a catalyst and also lowers the temperature for the decomposition of
- 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
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
2H 2 + O2 ¾¾Elec¾tric¾disc¾har¾ge ® 2H 2 O ;
N 2 + O2 ¾¾327¾3¾K ®
S + O2 ¾¾He¾at ® SO2 ; C + O2 ¾¾He¾at ® CO2
- Reaction with compounds : Dioxygen is an oxidising agent and it oxidises many compounds under
specific conditions. e.g. 4 HCl + O2 ¾¾CuC¾l ® 2H 2 O + 2Cl2 ;
4 NH 3 + 5O2 ¾¾107¾3¾K ® 4 NO + 6H 2 O
CS2 + 3O2 ¾¾He¾at ® 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
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
- Amphoteric oxides :
Al2 O3 , SnO2 etc.
- Neutral oxides :
H 2 O, CO, N 2O, NO
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,
|MgO basic|| Al2 O3
| SiO2 weakly
|P4 O10 acidic|| SO2 strongly
| Cl2O7 very
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 .
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.
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
2O 3 ¾¾475¾K ® 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 +
Therefore, ozone oxidises a number of non-metals and other reducing agents. e.g.
2Ag+ O3 ®
Ag 2 O + O2 ;
+ 3O3 ® SO3 + 3O2 ;
+ 4O3 ® PbSO4 + 4O2
Mercury is oxidised to mercurous oxide,
2Hg + O3 ®
Hg 2 O
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.
matter + O3 ® Oxidised coloured matter+ O2
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 ¾¾CC¾l4 ® H 2 C
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
and stable at room
dull yellow in colour, soluble in
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
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 ,
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
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
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 ;
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.
molecule has a bent structure with a O – S – O bond angle of 119o. Sulphur is
sp 2 hybridized.
- Sulphur trioxide (SO3) : It is formed by the oxidation of SO2 .
¾¾700¾K, 2¾at¾m. ® 2SO3
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 120o.
- 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
O = S – OH
O = S – OH
HO – S– S– OH
O = S — S = O
O = S – O– S = OH
HO = S — OOH
O = S – O – O – S = O
|diprotic, strong reducing|
|H 2 SO4 (+6)||Sulphuric acid||Stable diprotic, dehydrating|
|(Oil of vitriol)||agent|
|H 2S2O3 (–2 and +6)||
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|
|H 2 S2 O7 (+6)||Disulphuric acid||Strong oxidising agent|
|H 2 SO5 (+6)||
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 ;
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
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 ;
Pb(NO3 )2 + H 2 SO4 ® PbSO4 + 2HNO3
C12 H 22 O11 ¾¾Con¾c. H¾2SO¾4 ® 12C + 11H 2 O ; HCOOH ¾¾Con¾c.H¾2 SO¾4 ®
CO + H2O
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 ¾¾373¾K ®
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.
- 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
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
4 Na 2 S2 O3 ¾¾He¾at ® 3Na 2 SO4 +
Na 2 S5
- 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.
|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 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
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 +
3H 2O + 3F2 ® 6HF +
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 +
Br2 + H 2O ® HBr +
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
I 2 + KI ⇌
KI 2 or
I 2 + I – ⇌
- Reaction with hydrogen : Form covalent
H + F ¾¾–20¾0o¾C ® 2HF (very violent); H + Cl ¾¾Sunli¾g¾h¾t ® 2HCl
H 2 + Br2
¾¾He¾a¾t ® 2HBr ;
H 2 + I 2
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 103o, Cl2O is also V-shaped with bond angle 111o while ClO2 is angular with-bond angle 118o. 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|
|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
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
- HIO4 .
- Reaction with alkalies : 2F2 + 2NaOH ® 2NaF + OF2 + H 2 O
; 2F + 4 NaOH ® 4 NaF + O2 + 2H 2 O
Halogen other than fluorine (Cl 2 ,Br2 , I 2 ) react with NaOH as follows,
X 2 (g) +
¾¾15o¾C ® X – + OX – + H O ;
X (g ) + 6OH – ¾¾70o¾C ® 5X – +
+ 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.
BrF, BrCl, ICl
| ClF3 , BrF3
IF3 , ICl3
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,
⇌ 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
|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,
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 ¾¾Elec¾tric¾ity ® 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 ¾¾MnO¾2 ® 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.
|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
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
|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%.
- 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
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
in gas phase and a orange yellow solid of the
composition Xe PtF6 was obtained,
Xe(g) + PtF6(g) ® Xe + [PtF6
Some important stable compounds of Xe are,
|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 ¾¾Ni, 6¾73¾K ® XeF2
Structure : XeF2 has trigonal bipyramid geometry due to
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.
is also formed when an electric discharge is
passed through a mixture of xenon and excess of fluorine,
Xe + 2F2 ¾¾Ni, 6¾73¾K ® XeF4
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.
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 ¾¾193¾K ® 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
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
pp – dp
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
pp – dp
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
and xenon dioxydifluoride (XeO2F2).
- Xenon oxydifluoride (XeOF2) is formed by partial hydrolysis of XeF4 at 193 K,
XeF4 + H 2 O ¾¾193¾K ® 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
pp – dp
- 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
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
XeO3 + XeOF4 ¾¾195¾K ® 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
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