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OXYGEN FAMILY

             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.
 

(1) Electronic configuration


Physical properties
(1) Physical state : Oxygen is gas while all other are solids.
(2) Atomic radii : Down the group atomic radii increases because the increase in the number of inner shells overweighs the increase in nuclear charge.
(3) Ionisation energy : Down the group the ionisation energy decrease due to increase in their atomic radii and shielding effect.
(4) Electronegativity : Down the group electronegativity decreases due to increase in atomic size.
(5) Electron affinity : Element of this group have high electron affinity, electron affinity decreases down the group.
(6) Non – metallic and metallic character : These have very little metallic character because of their higher ionisation energies.
(7) Nature of bonding : Compound of oxygen with non metals are predominantly covalent. S, Se, and Te because of low electronegativities show more covalent character.
(8) Melting and boiling points : The melting point and boiling points increases on moving down the group.

(9) Catenation : Oxygen has some but sulphur has greater tendency for catenation.
H - O - O - H,        H - S - S - H.
      (H2O2)                          (H2S2)
H - S - S - S - H,          H - S - S - S - S - H
      (H2S3)                         (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 – α and β (both metallic)
(11) Oxidation states : Oxygen shows – 2, + 2 and –1 oxidation states. Other elements show +2 ,+4 and +6 oxidation states.
Chemical properties
(1) Hydrides : The elements of this group form hydrides such as H2O, H2S, H2Se, H2Te an H2Po. Following are their characteristics.

(i) Physical states : Water is colourless and odourless while hydrides of the rest of the elements of this group are colourless, unpleasant smelling poisonous gases.
(ii) Volatile nature : Volatility increases from H2O  to H2S  and then decreases. The low volatility and abnormally high boiling point of water is due to the association of water molecules on account of hydrogen bonding because of strongly electronegative oxygen atom linked to hydrogen atom. thus, water is liquid while H2S  and other hydrides are gases under normal condition of temperature and pressure.
(iii) Acidic character : The hydrides of this group behave as weak diprotic acids in aqueous solution, the acidic character increasing from H2S  to H2Te when H2O is neutral.
(iv) Thermal stability : The thermal stability decreases from H2O to H2Po because the size of the central atom (from O to Po ) increases resulting in longer and weaker M − H bond consequently the bond strength decreases. This results in the decrease of the thermal stability.
(v) Reducing character : The reducing power of the hydrides increases from H2O to  H2Po due to the decreasing bond strength from H2O to H2Po.

(vi) Bond angle : All these hydrides are angular molecules and the bond angle H − X − H ( X is O , S , Se, Te ) decreases from H2O to H2Te .
         Increasing order of reducing power of hydrides :
H2O < H2S < H2Se < H2Te
Increasing order of bond angles in hydrides :
H2Te < H2Se < H2S < H2O
The order of stability of hydrides :
H2O > H2S > H2Se > HTe
The order of increasing acidic nature of hydrides :
H2O < H2S < H2Se < H2Te
(2) Oxides: These elements form monoxides (MO), dioxides (MO2) and trioxides (MO3).
(i) Dioxides: Sulphur, selenium and tellurium burn in air to form SO2, SeO2 and TeO2. The dioxide molecules contain pπ  − pπ  bonds which become weaker with increase in atomic number because of the increase in the bond length.
(a) Sulphur dioxide, SO2 is a gas at room temperature and exists as individual molecules even in the solid state. Its molecule has bent structure and is a  resonance hybrid of the following canonical structures.


             SO2 is acidic in nature and also called the anhydride of sulphurous acid. It can act as reducing and oxidising agent. SO2 also acts as a beleaching agent in the presence of moisture, but in contrast to Cl2 , its bleaching action is temporary.
SO2 + 2H2O  H2SO4 + 2[H]
Colouring matter +2 [H]  Colourless compound Hence, SO2 bleaches due to reduction and the bleaching action is temporary.
(b) Selenium dioxide, SeO2 is a solid with polymeric zig‐zag structure at room temperature however it exist as discrete molecules in the gaseous phase.


                            
(c) Tellurium dioxide, TeO2  is also a solid with polymeric zig‐zag structure at room temperature very similar to that of selenium dioxide.

(ii) Trioxides : Sulphur, selenium and tellurium can form trioxides also.
(a) Sulphur trioxide, SO3 : In the gaseous state monomeric SO3 has a planar structure with S − O bond distance of 143 pm and O − S − O  bond angle of 120o. SO3 molecule is a resonance hybrid of following structures.

           In the solid phase sulphur trioxide polymerises to cyclic trimer or to a stable linear chain structure. SO3 is the anhydride of H2SO4. It is acidic in nature and acts as oxidising agent.
(b) Selenium trioxide, SeO3 : it is a solid substance which exists as a cyclic tetramer, however in the vapour phase it exists as a monomer
(c) Tellurium trioxide, TeO3 : It is a solid at room temperature existing as a polymer.
The increasing order of acidic nature of oxides is TeO3 < SeO3 < SO3 .

(3) Oxyacids :
H2SO3, H2SO4, H2S2O3, H2SO5, H2S2O8, H2S2O7, H2S2O6
(4) Halides : Oxygen : OF2, Cl2O, Br2O
Sulphur : S2F2, S2Cl2, SF2, SCl2, SBr2, SF4, SCl4 and SF6
Selenium and tellurium : SeF6 and TeF6

 

Anamolous Behaviour of Oxygen
Oxygen is the first member of the group 16 family and differs from the other members of the family because of
(1) Its small size
(2) Its high electronegativity
(3) Its high ionisation energy
(4) Absence of d -orbitals in the valence shell
It differs from the other members of the family as follows
(1) Elemental state : Oxygen is a diatomic gas while others are octa-atomic solids with eight membered puckered ring structure.

(2) Oxidation states : Oxygen shows O.S. of – 2 in most of its compounds. It also shows an O. S. of +2 in F2O and –1 in  H2O2 or other peroxides. It cannot show O.S. beyond 2. Other elements show oxidation states of +2, +4 and +6 because these elements have vacant d -orbitals so that their valence shell can expand.
(3) Hydrogen-bonding : Oxygen atom is very small and has quite high nuclear charge. therefore, it has high value of electronegativity and is able to form H -bonds. the other elements, because of their large size, cannot form H -bonds. As a result, H2O is liquid while H2S  is a gas and H2Se etc., are solids.
(4) Maximum covalency : Oxygen has a maxium covalency of two while other elements can show a maximum covalency of six. This is because these elements have vacant d -orbitals while oxygen has not.
(5) Types of compounds : The compounds of oxygen are mainly ionic and polar covalent due to high electronegativity of oxygen while those of others are not.
(6) Magnetic character : Oxygen is paramagnetic while others are not.

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 :
                8O16                                   8O17                            8O18
(Abundance : 99.76%)       (Abundance : 0.037%)      (Abundance : 0.204%)
Out of the three isotopes, 8O18 is radioactive.
Occurrence : In free state, it occurs in air and constitutes 21% by volume of air.
Preparation of Dioxygen : Oxygen is prepared by the following methods.
(1) By the decomposition of oxygen rich compounds : e.g.

(2) By heating dioxides, Peroxides and higher oxides : e.g.

(3) Laboratory Method : In the laboratory, O2 is prepared by thermal decomposition of potassium chlorate.

           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.
(4) O2 can also be prepared by the action of water on sodium peroxide as,
2Na2O2 + 2H2O 4 NaOH + O2
(5) Industrial preparation : The main sources for the industrial preparation of dioxygen are air and water.
(i) From air : O2 is prepared by fractional distillation of air. During this process, N2 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.
(ii) From water : O2 can also be obtained by the electrolysis of water containing a small amount of acid or alkali, 

Physical properties of O2 : It is a colourless, tasteless and odourless gas. It is slightly soluble in water and its solubility is about litre per 30 cm3 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 Acidic to basic character increases and its bond dissociation energy is very high.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.

(1) Action with litmus : Like dihydrogen, it is also neutral and has no action on blue or red litmus.
(2) Reaction with metals : Active metals like Na, Ca react at room temp. to form their respective oxides.
4Na + O2  2Na2O;   2Ca + O2  2CaO
It reacts with Fe, Al, Cu etc. metals at high temperature
4Al + 3O2  2Al2O3;   4 Fe + 3O2 2Fe2O3
(3) Action with Non‐metals : It form oxides.


         

(4) Reaction with compounds : Dioxygen is an oxidising agent and it oxidises many compounds under specific conditions. e.g.

Uses of dioxygen
(1) It is used in the oxy-hydrogen or oxyacetylene torches which are used for welding and cutting of metals.
(2) It is used as an oxidising and bleaching agent,
(3) Liquid O2 is used as rocket fuel.
(4) It is used in metallurgical processes to remove the impurities of metals by oxidation.

 

Compounds of Oxygen
(1) 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,

(i) Basic oxides : Alkali, alkaline earth and transition metals form basic oxides ‐ Na2O, MgO, Fe2O3 etc. their relative basic character decreases in the order : alkali metal oxides > alkaline earth metal oxides > transition metal oxides.
(ii) Acidic oxides : Non‐metal oxides are generally acidic
‐ CO2, SO2, SO3, NO2, N2O5, P4O10, Cl2O7 etc.
(iii) Amphoteric oxides : Al2O3, SnO2 etc.
(iv) Neutral oxides : H2O, CO, N2O, 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,

                       

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:


                        
                                
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 , H2O, CaO, Li2O, Al2O3 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 O22−  ion having oxidation number of oxygen as –1. For example,
H2O2, Na2O2, BaO2, pbO2 etc.

Superoxides : These contains O2 ion having oxidation number of oxygen as –1/2. For example, KO2, PbO2 etc.
Suboxides : These oxides contain less oxygen than expected from the normal valency. For example, N2O.
Mixed oxides : These oxides are made up of two simple oxides. For example, red lead Pb3O4(PbO + PbO2). magnetic oxide of iron, Fe3O4(FeO + Fe2O3) and mixed oxide of manganese, Mn3O4(MnO2 + 2MnO).

 

Ozone or trioxygen
              Ozone is an allotrope of oxygen. It is present in the upper atmosphere, where it is formed by the action of U. V. radiations on O2 ,

O3 protects us from the harmful U. V. radiations which causes skin cancer. Now a days, ozone layer in the atmosphere 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.


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 (≈273 K ) must be maintained. (d) The electric discharge must be sparkless.
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,
(1) Decomposition : Pure ozone decomposes on heating above 475 K to form O2 gas.


                      

(2) Oxidising agent : Ozone is one of the most powerful oxidising agent with the liberation of dioxygen. In fact, ozone is a stronger oxidising agent than molecular oxygen because ozone has higher energy content and decomposes to give atomic oxygen as:


                       
Therefore, ozone oxidises a number of nonmetals and other reducing agents. e.g.


            
Mercury is oxidised to 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.
(3) Bleaching agent : Due to the oxidising action of ozone, it acts as a mild bleaching agent as well as a sterilizing agent. It acts as a bleaching agent for vegetable colouring matter.
For example, ozone bleaches indigo, ivory, litmus, delicate fabrics etc.

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


                      
Structure of O3 : The structure 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
(1) O3 is used for disinfecting water for drinking purposes because ozone has germicidal properties.

(2) It is used for purifying air of crowded places such as cinemas, under ground railway, auditoriums, tunnels, mines etc.
(3) 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 0.05%. It occurs in the free state as well as in combined state. Sulphur occurs mainly as sulphides and sulphates. eg.


               

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,
(1) Rhombic or octahedral or α‐sulphur : It is a bright yellow solid, soluble in CS2 and stable at room temp. All other varieties of sulphur gradually change into this form on standing.
(2) Monoclinic sulphur or prismatic or β‐ 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 out. It is 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.
(3) Plastic or amorphous or γ ‐sulphur : It is a super cooled liquid insoluble in CS2, soft and amorphous. It consists of long zig‐zag chains of S - atoms.


                           

(4) Colloidal or δ ‐sulphur : It is prepared by passing H2S through a solution of an oxidizing agent or water or by treating sodium thiosulphate with dil. HCl.
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 H2SO4 to SO2. With NaOH solution on heating,
                           S8 + 12NaOH   4Na2S + 2Na2S2O3 + 6H2O.
It gives sodium sulphide and sodium thiosulphate, with excess of sulphur,
2Na2S + S8  2Na2S5.
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,H2SO4, CS and dyes, sulpha drugs and ointment for curing skin diseases and in the vulcanization of rubber.
Compounds of Sulphur
(1) Hydrogen Sulphide: It is prepared in the laboratory by the action of dill. H2SO4 on ferrous sulphide in kipp's apparatus, Fes + H2SO4  FeSO4 + H2S. 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,

The solubility of sulphides can be controlled by the H+ ions concentration and therefore, H2S finds extensive use in qualitative analysis of cation radicals.


                                              
(2) Halides of sulphur : Two important halides of sulphur are SF4 and SF6.
(i) Sulphur tetrafluoride : SF4 is formed by the reaction of sulphur with CoF3.
            S + 4CoF3  SF4 + 4CoF2
 It is a colour gas which is quite reactive. It is hydrolysed with water.

SF4 + 2HO  SO2 + 4HF
It is used for fluorinating inorganic and organic compounds.
Structure : It has see‐saw structure with  sp3d ‐ hybrdization and is derived from triogonal bipyramid geometry in which an equatorial position is occupied by a lone pair of electrons.


                         
(ii) Sulphur hexafluoride : SF6 is prepared by burning sulphur in a stream of fluorine. OF6 is not known though sulphur forms SF6. This is because oxygen has no d ‐orbitals in its valence shell.
            SF6 is a colourless gas. It is extremely inert substance even at red heat. It does not react with water. on account of its chemical inertness and dielectric strength, it is used as an insulator in high voltage generators and switch‐gears.

Structure : It has an octahedral structure with sp3d2 ‐ hybridisation around the central sulphur atom.
Therefore, all  S − F  bond distances are equal in its structure.


                                          
(3) Oxides of sulphur : Sulphur forms several oxides of which sulphur dioxide (SO2) and sulphur trioxide (SO3) are most important.
(i) Sulphur dioxide (SO2) : It is prepared by burning sulphur or iron pyrites in air.
S8 + 8O2  8SO2;
4FeS2 + 11O2  2Fe2O3 + 8SO2
In laboratory, it is prepared by heating copper turnings with conc. H2SO4
Cu + 2H2SO4  CuSO4 + SO2 + 2H2O

It is colourless gas with irritating and suffocating smell.
SO2 molecule has a bent structure with a O – S – O bond angle of 119º. Sulphur is sp2  hybridized.


                                                  

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


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


                       
(4) Oxyacids of sulphur: Sulphur forms many oxyacids. Some of these are,

Sulphuric acid (H2SO4) : H2SO4 is a very stable oxyacid of sulphur. 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 catalyst.
S + O2  SO2 or 4FeS2 + 11O2  2Fe2O3 + 8SO2 
              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, (a) High concentration of SO2 and O2. (b) Low temperature of 673 to 723 K, (c) 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.

 Contact process is preferred over lead chamber process (gives 98% pure H2SO4) since it gives H2SO4 of greater purity (100%).
Structure : H2SO4 is a covalent molecule with sulphur in a +6 oxidation state. The two oxygen atoms are linked to sulphur by double bonds while the other two oxygen atoms.
Are linked by single covalent bonds. Thus it has tetrahedral structure. Infact, sulphuric acid has an associated structure due to the presence of hydrogen bonds. As a result, it is a dense and viscous liquid and has a high boiling point of 590K


                  

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.
(i) H2SO4 is a strong dibasic acid. It neutralizes alkalies, liberates CO2 from carbonates and bicarbonates.
(ii) It reacts with more electropositive (than hydrogen) metals to evolve H2 and produces SO2 on heating with less electropositive metals than hydrogen .eg.,
H2SO4 + 2KOH  K2SO4 + 2H2O;


Cu + 2H2SO4   CuSO4 + SO2 + 2H2O

(iii) It is a strong Oxidizing agent and Oxidises as follows,
H2SO4  H2O + SO2 + O
C + H2SO4   2SO + CO + 2H2O
S + 2H2SO4   3SO2 + 2H2O
P4 + 10H2SO4   4H2PO4 + 10SO2 + 4H2O
2HBr + H2SO4   Br2 + 2H2O + SO2
2HI + H2SO4  2H2O + I2 + 2SO2
(iv) It reacts with number of salts. It liberates HCl from chlorides, H2S from sulphides, HNO3 from nitrates.
(v) 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 ethylene.
(vi) It is also a good sulphonating agent and used for sulphonation of aromatic compounds. eg.,


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 : Na2S2O3 . 5H2O It is manufactured by saturating a solution of sodium carbonate with SO2 which gives a solution of sodium sulphite,
Na2CO3 + SO+ H2O  Na2SO3 + CO2 + H2O
The resulting solution is boiled with powdered sulphur as, 
The solution is then cooled to get crystals of sodium thiosulphate.
Physical properties : (i) Sodium thiosulphate is a colourless crystalline solid. In the hydrated form, it is called hypo. (ii) It melts at 320 K and loses its water molecules of crystallization on heating to 490K.

Chemical properties
(i) Action with halogens : It reacts with halogens as,
(a) Chlorine water oxidizes sodium thiosulphate to sodium sulphate and sulphur is precipitated,
Na2S2O3 + Cl2 + H2O  2HCl + 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.
(b) Bromine water also oxidizes sodium thiosulphate to sodium sulphate and sulphur,
Na2S2O3 + Br2 + H2  Na2SO4 + 2HBr + S
(c) With iodine it forms a soluble compound called sodium tetrathionate,

Therefore, hypo is commonly used to remove iodine stains from the clothes.
(ii) Action of heat : Upon heating, sodium thiosulphate decomposes to form sodium sulphate and sodium pentasulphide,


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

Na2S2O3 + 2HCl 2NaCl + SO2 + H2O + S
(iv) Action with silver halides : Sodium thiosulphate forms soluble complex when treated with silver chloride or silver bromide,

This property of hypo is made use in photography.
Uses of sodium thiosulphate
(i) It is largely used in photography as a fixing agent.
(ii) It is used as a preservative for fruit products such as jams and squashes.
(iii) It is used as an antichlor in bleaching.
(iv) It is used as a volumetric agent for the estimation of iodine.
(v) It is used in medicine.

Posted Date : 19-02-2021

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