1. Group 15 elements

Chapter 7

p-block (nitrogen and oxygen family)

Introduction :

Group 13 to 18 of the periodic table of elements constitute the p–block. The p–block contains metals, metalloids as well as non–metals.

The  p–block elements have general valence shell electronic configuration ns2 np1–6.

The first member of each group from 13–17 of the p–block elements differ in many respects from the other members of their respective groups because of small size, high electronegativity and absence of d–orbitals.

The first member of a group also has greater ability to form pp–pp multiple bonds to itself (e.g. C=C, CºC, NºN) and to element of second row (e.g C=O, C=N, CºN, N=O) compared to the other members of the same group.

 The highest oxidation of p–block element is equal to the group number minus 10. Moving down the group, the oxidation state two less than the highest group oxidation state becomes more stable in groups 13 to 16 due to inert pair effect (reluctance of s-subshell electrons to participate in chemical bonding)

Group 15 Elements : The Nitrogen family

Group 15 includes nitrogen phosphorus, arsenic, antimony and bismuth. As we go down the group, there is a shift from non-metallic to metallic through metalloidic character. Nitrogen and phosphorus are non-metal, arsenic and antimony metalloid and bismuth is a typical metal.

Electronic Configuration : The valence shell electronic configuration of  these element is ns2 np3. The s orbital in these element is completely filled and p orbitals are half- filled, making their electronic configuration extra stable.

Atomic and Ionic Radii : Covalent and ionic (in a particular state) radii increase in size down the group. There is a considerable increase in covalent radius from N to P. However, from As to Bi only a small increase in covalent radius is observed. This is due to the presence of completely filled d and / or f orbitals in heavier members.

Ionisation Enthalpy : Ionisation enthalpy decreases down the group due to gradual increase in atomic size. Because of the extra stable half- filled p-orbital electronic configuration and smaller size, the ionisation enthalpy of the group 15 element is much greater than of group 14 elements in the corresponding periods. The order of successive ionisation enthalpies, as expected is DiH1 < DiH2 < DiH3 

Electronegativity : The electronegativity value, in general, decreases down the group with increasing atomic size. However, amongst the heavier elements, the difference is not that much pronounced.

Physical Properties : All the elements of this group are polyatomic. Dinitrogen is a diatomic gas while all others are solids. Metallic character increases down the group. Nitrogen and phosphorus are non–metals, arsenic and antimony metalloids and bismuth is a metal. This is due to decrease in ionisation enthalpy and increase in atomic size. The boiling points, in general, increase from top to bottom in the group but the melting point increases upto arsenic and then decreases upto bismuth. Except nitrogen, all the elements show allotropy.

Atomic & physical properties


Chemical Properties :   

Oxidation States and trends in a chemical reactivity :

The common oxidation states of these elements are –3, +3 and +5. The tendency to exhibit –3 oxidation state decreases down the group , bismuth hardly forms any compound in –3 oxidation state. The stability of +5 oxidation state decreases down the group. The only well characterised Bi (V) compound is BiF5 .The stability of +5 oxidation state decreases and that of +3 state increases (due to inert pair effect) down the group.

 Bi3+  >  Sb3+  >  As3+   ;   Bi5+  <  Sb5+  <  As5+

Nitrogen exhibits +1, +2, +4 oxidation states also when it reacts with oxygen. Phosphorus also shows +1 and +4 oxidation states in some oxoacids.

In the case of nitrogen, all oxidation states from +1 to +4 tend to disproportionate in acid solution.

For example,

3 HNO2  ® HNO3 + H2O + 2 NO

Similarly, in case of phosphorus nearly all intermediate oxidation states disproportionate into +5 and –3 both in alkali and acid. However +3 oxidation state in case of arsenic , antimony and bismuth become increasingly stable with respect to disproportionation.

Nitrogen is restricted to a maximum covalency of 4 since only four (one s and three p) orbitals are available for bonding. The heavier elements have vacant d orbitals in the outermost shell which can be used for bonding (covalency) and hence , expand their covalence as in PF6 .

Anomalous properties of nitrogen :

Nitrogen differs from the rest of the members of this group due to its smaller size, high electronegativity, high ionisation enthalpy and non–availability of d orbitals. Nitrogen has unique ability to form pp–pp multiple bonds with itself and with other elements having small size and high electronegativity (e.g., C, O). Heavier elements of this group do not form pp–pp bonds as their atomic orbitals are so large and diffuse that they cannot have effective overlapping. Thus, nitrogen exists as a diatomic molecule with a triple bond (one s and two p) between the two atoms. Consequently, its bond enthalpy (941.1 kJ mol–1) is very high. On the contrary, phosphorus, arsenic and antimony form metallic bonds in elemental state. However, the single N–N bond is weaker than the single P–P bond because of high interelectronic repulsion of the non–bonding electrons, owing to the small bond length. As a result the catenation tendency is weaker in nitrogen. Another factor which affects the chemistry of nitrogen is the absence of d orbitals in its valence shell. Besides restricting its covalency to four, nitrogen cannot form dp–pp bonds as the heavier elements can e.g., R3P=O or R3P=CH2 (R = alkyl group). Phosphorus and arsenic can form dp–pp bond also with transition metals when their compounds like P(C2H5)3 and As(C6H5)3 act as ligands.

(i) Reactivity towards hydrogen : All the elements of Group 15 form hydrides of the type EH3 where E=N, P, As, Sb or Bi. Some of the properties of these hydrides are shown in Table. The hydrides show regular gradation in their properties. The stability of hydrides decreases from NH3 to BiH3 which can be observed from their bond dissociation enthalpy. Consequently, the reducing character of the hydrides increases. Ammonia is only a mild reducing agent while BiH3 is the strongest reducing agent amongst all the hydrides. Basicity also decreases in the order NH3 > PH3 > AsH3 > SbH3 ³ BiH3 .

Properties of Hydrides of Group 15 Elements

(ii) Reactivity towards oxygen : All these elements form two types of oxides : E2O3 and E2O5 . The oxide in the higher oxidation state of the element is more acidic than that of lower oxidation state. Their acidic character decreases down the group. The oxides of the type E2O3 of nitrogen and phosphorus are purely acidic , that of arsenic and antimony amphoteric and those of bismuth is predominantly basic.

(iii) Reactivity towards halogens : These elements react to form two series of halides : EX3 and EX5 . Nitrogen  does not form pentahalide due to non – availability of the d-orbitals in its valence shell. Pentahalides are more covalent than trihalides. All the trihalides of these elements except those of nitrogen are stable. In case of nitrogen, only NF3 is known to be stable. Trihalides except BiF3 are predominantly covalent in nature. Halides are hydrolysed in water forming oxyacids or oxychlorides.

 PCl3 + H2® H3PO3 + HCl

 SbCl3 + H2®  SbOCl¯ (orange) + 2HCl

BiCl3 + H2®  BiOCl¯ (white) + 2HCl

(iv)   Reactivity towards metals : These elements react with metals to form their binary compounds exhibiting –3 oxidation state , such as , Ca3N2 (calcium nitride) Ca3P2 (calcium phosphide) , Na3As2 (sodium arsenide), Zn3Sb2 (zinc antimonide) and Mg3Bi2 (magnesium bismuthide).

2. Dinitrogen


Preparation :

(i)     Laboratory method of preparation : 

NH4Cl (aq) + NaNO2 ® (aq)  N2 (g) + 2 H2O (1) + NaCl (aq)

Small amounts of NO and HNO3 are also formed in this reaction ; these impurities can be removed by passing the gas through aqueous suplhuric acid containing potassium dichromate.

N2 is collected by the downward displacement of water.

(ii)    By heating ammonium dichromate :  (NH4)2Cr2O7  N2 ­ + 4H2O + Cr2O3

(iii) Very pure nitrogen   ;  Ba(N3)2  Ba + 3N2 

Sodium azide also gives N2 on heating.     

(iv) Industrial method of preparation :

From liquefied air by fractional distillation : The boiling point of N2 is –196oC and that of oxygen is –183oC and hence they can be separated by distillation using fractional column.

Properties :

(i) N2 is a colourless, odourless, tastless, non-toxic gas having very low solubility in water (23.2 cm3 per litre water at 273 K and 1 bar pressure). It has two stable isotopes : 14N and 15N. It is neither combustible nor a supporter of combustion.

(ii)   Li, Mg and Al on heating with N2 form corresponding nitrides.

6Li + N2  2Li3N ; 3Mg + N2  Mg3N2 ; 2Al + N2  2AlN

(iii) Reaction with H2 : At 200 atm and 773 K, and in the presence of iron catalyst and molybdenum promoter, N2 combines with H2 reversibly to form ammonia. The process is called Haber’s Process and is the industrial method of manufacturing ammonia. The reaction is exothermic.

N2 + 3H2  ® 2NH3

(iv)Reaction with oxygen: When air free from CO2 and moisture is passed over an electric arc at about 2000 K, nitric oxide is formed. This reaction is endothermic.

  N2 + O2 ® 2NO

Uses :

1.  For providing an inert atmosphere during many industrial processes where presence of air or O2 is to be avoided (e.g., in iron and steel industry, inert diluent for reactive chemicals).

2. For manufacture of NH3 by the Haber’s process.

3. For manufacture of HNO3 by the Birkeland-Eyde process.

4. For manufacture of industrial chemicals containing nitrogen like calcium cyanamide.

5. Liquid dinitrogen is used as a refrigerant to preserve biological materials, food items and in cryosurgery.

3. Ammonia


Preparation :

(i)   Ammonia is present in small quantities in air and soil where it is formed by the decay of nitrogenous orgainc matter e.g., urea.

NH2CONH2 + 2 H2O ®  (NH4)2CO3   2 NH3 + H2O  + CO2

(ii)  On a small scale ammonia is obtaned from ammonia salts which decompose when treated with caustic soda or lime.

2 NH4Cl + Ca (OH)2 ® 2 NH3 + 2 H2O + CaCl2          

(NH4)2 SO4 + 2 NaOH ® 2 NH3 + 2 H2O + Na2SO4

(iii)   On a large scale, ammonia is manufactured by Haber’s process.

  N2 (g) + 3 H2 (g) ® 2 NH3 (g) ;      Df H = – 46.1 kJ mol–1

In accordance with Le Chatelier’s principle, high pressure would favour the formation of ammonia. The optimum conditions for the production of ammonia are a pressure of 200 × 105 Pa (about 200 atm) , a temperature of ~ 700 K and the use of a catalyst such as iron oxide with small amounts of K2O and Al2O3 to increase the rate of attainment of equilibrium.

(iv)   By hydrolysis of metal nitrides like AIN or Mg3N2         

AIN + NaOH + H2O ® NaAIO2 + NH3 

For drying, dehydrating agents like H2SO4 , P2O5 or CaCl2 can not be used as these react with NH3.

2NH3 + H2SO4 ® (NH4)2SO4 ; 6NH3 + P2O5 + 3H2O ® 2(NH4)3PO4 

 CaCl2 + 8NH3  ® CaCl2 8NH3 (Adduct)

So quicklime (CaO) is used for drying of NH3 .

CaO + H2O ® Ca(OH)2

Properties :

Physical properties : Ammonia is a colourless gas with a pungent odour. Its freezing and boiling points are 198.4 and 239.7 K respectively. In the solid and liquid states , it is associated through hydrogen bonds and that accounts for its higher melting and boiling points than expected on the basis of its molecular mass. The ammonia molecule is trigonal pyramidal with the nitrogen atom at the apex. It has three bond pairs and one lone pair of electrons as shown in the structure. 

Ammonia gas is highly soluble in water. Its aqueous solution is weakly basic due to the formation of OH ions.  NH3 (g) + H2O (1)  NH4+ (aq) + OH (aq)

Chemical properties :

(i)  It forms ammonium salts with acids , e.g., NH4Cl , (NH4)2 SO4 , etc. As a weak base , it precipitates the hydroxides of many metals from their salt solutions. For example ,

 2 FeCl3 (aq) + 3 NH4OH (aq) ® Fe2O3 . xH2O (s) + 3 NH4Cl (aq)

                                                                                (brown ppt)

(ii) The presence of lone pair of electrons on the nitrogen atoms of the ammonia molecule makes it a Lewis base. It donates the electrons pair and forms linkage with metal ions forming complex compounds.

Cu2+ (aq) + 4 NH3 (aq)   [Cu(NH3)4]2+ (aq)

(blue)                                            (deep blue)

Ag+ (aq) + Cl–1 (aq)  AgCl (s)

(colourless)                      (white ppt)

 AgCl (s) + 2 NH3 (aq)  ® [Ag (NH3)2]Cl (aq)

(white ppt)                          (colourless)

Test of ammonia/ammonium salts :

When NH3 gas is passed into the colourless solution of Nessler’s reagent a brown precipitate or coloration is formed. This is a test for NH3 gas.

2K2HgI4 + 3KOH + NH® H2N·HgO·HgI (brown) + 7KI + 2H2O

Uses :

1. Liquid ammonia is used as a refrigerant.

2. For the production of ammonium fertilizers such as ammonium sulphate, ammonium phosphate, ammonium nitrate, urea etc.

3. For removing grease because NH4OH dissolves grease.

4. For manufacture of HNO3 by the Ostwald process.

5. As a laboratory reagent.

6. In the production of artificial rayon, silk, nylon etc

4. Oxides of nitrogen

Oxides of nitrogen

Nitrogen forms a number of oxides, N2O, NO, N2O3, NO2 or N2O4 and N2O5, and also very unstable NO3 and N2O6. All these oxides of nitrogen exhibit pπ-pπ multiple bonding between nitrogen and oxygen.

5. Nitric acid


Preparation :

(i)    In the laboratory, nitric acid is prepared by heating KNO3 or NaNO3 and concentrated H2SO4 in a glass retort.

NaNO3 + H2SO4  ® NaHSO4 + HNO3 

(ii)   On a large scale it is prepared mainly by Ostwald’s process.

This method is based upon catalytic oxidation of NH3 by atmospheric oxygen.

4 NH3 (g) + 5 O2 (g) (from air)  4 NO (g) + 6 H2O (g)

  Nitric oxide thus formed combines with oxygen giving NO2.

 2 NO (g) + O2 (g)  2 NO2 (g)

Nitrogen dioxide so formed, dissolves in water to give HNO3.

3 NO2 (g) + H2O (l) ® 2 HNO3 (aq) + NO (g)

NO thus formed is recycled and the aqueous HNO3 can be  concentrated by distillation upto ~ 68% by mass. Further concentration to 98% can be achieved by dehydration with concentrated H2SO4.

Properties :

Physical properties :

It is a colourless liquid. Freezing point is 231.4 K and boiling point is 355.6 K. Laboratory grade nitric acid contains ~ 68% of the HNO3 by mass and has a specific gravity of 1.504.

In the gaseous state, HNO3 exists as a planar molecule.

In aqueous solution, nitric acid behaves as a strong acid giving hydronium and nitrate ions.

HNO3 (aq) + H2O (l® H3O+ (aq) + NO3 (aq)

(i) Concentrated nitric acid is a strong oxidising agent and attacks most metals except noble metals such as gold and platinum. The products of oxidation depend upon the concentration of the acid , temperature and the nature of the material undergoing oxidation.

Some metals (e.g., Fe, Cr , Al) do not dissolve in concentrated nitric acid because of the formation of a passive film of oxide on the surface.

1.  4Zn + 10HNO3 (dilute) ® 4Zn(NO3)2 + N2O  + 5H2O

 Zn + 4HNO3 (concentrated) ® Zn(NO3)2 + 2NO2 + 2H2O

2.  3Cu + 8HNO3 (dilute) ® 2NO + Cu(NO3)2 + 4H2O

Cu + 4HNO3 (concentrated)  ® 2NO2 + Cu(NO3)2 + 2H2O

(ii)   Concentrated nitric acid also oxidises non–metals and their compounds. Iodine is oxidised to iodic acid , carbon to carbon dioxide , sulphur to H2SO4 and phosphorus to phosphoric acid.

I2 + 10 HNO3 ® 2 HIO3 + 10 NO2 + 4 H2O

C + 4 HNO3 ® CO2 + 2 H2O + 4 NO2 

S8 + 48 HNO3 (concentrated) ® 8 H2SO4 + 48 NO2 + 16 H2O

P4 + 20 HNO3 (concentrated) ® 4 H3PO4 + 20 NO2 + 4 H2O

Brown Ring Test :

The familiar brown ring test for nitrates depends on the ability of Fe2+ to reduce nitrates to nitric oxide, which reacts with Fe2+ to form a brown coloured complex. The test is usually carried out by adding dilute ferrous sulphate solution to an aqueous solution containing nitrate ion , and then carefully adding concentrated sulphuric acid along the sides of the test tube. A brown ring at the interface between the solution and sulphuric acid layers indicate the presence of nitrate ion in solution.

 NO3 + 3 Fe2+ + 4H+ ® NO + 3Fe3+ + 2 H2O

[Fe (H2O)6]2+ + NO  ® [Fe (H2O)5 (NO)]2+ + H2O  

Uses :

The major use of nitric acid is in the manufacture of ammonium nitrate for fertilizers and other nitrates for use in explosives and pyrotechnics. It is also used for the preparation of nitroglycerin, trinitrotoluene and other organic nitro compounds. Other major uses are in the pickling of stainless steel, etching of metals and as an oxidiser in rocket fuels.

6. Phosphorus - Allotropic forms


It occurs in nature in the form of stable phosphates. (Animal bones also contain calcium phosphate (58 %)). The important minerals are:

(i) Phosphorite, Ca3(PO4)2 

(ii) Chloraptite, Ca3(PO4)2CaCl2

(iii) Fluoraptite, Ca3(PO4)2CaF2  

(iv) Vivianite, Fe3(PO4)2·8H2O (v) Redonda phosphate, AlPO4

Phosphorus Allotropic Forms :

 Phosphorus is found in many allotropic forms , the important one being white , red and black.

White phosphorus :

Preparation :
2Ca3(PO4)2 (f`rom bone-ash) + 10C + 6SiO2  6CaSiO3 + 10CO + P4(s) (electric furnace method)

It is a translucent white waxy solid. It is poisonous , insoluble in water but soluble in carbon disulphide. Molecular formula is P4. Ignition temperature is around 30ºC.

When exposed to air it undergoes oxidation which gradually raises it temperature and ultimately catches fire when the temperature exceeds 30ºC. That is why it is kept in water. White phosphorus is less stable and therefore , more reactive than the other solid phases under normal conditions because of angular strain in the P4 molecule where the angles are only 600.

It glows in dark due to slow oxidation. This property is called phosphorescence (chemiluminescence).

P4 + 5O2  ® P4O10                                         

 It consists of discrete tetrahedral P4 molecule as shown in Fig.                      

Red phosphorus is obtained by heating white phosphorus at 573 K in an inert atmosphere of CO2 or coal gas for several days. This red phosphorous may still contain some white phosphorus which is removed by boiling the mixture with NaOH when white phosphorus is converted in to PH3 gas but red phosphorus remains inert.

P4 + 3NaOH + 3H2® PH3(g) + 3NaH2PO2 

When red phosphours is heated under high pressure, a series of phase of black phosphorus are formed. Red phosphorus possesses iron grey lustre. It is odourless , non – poisonous and insouble in water as well as in carbon disulphide. Chemically, red phosphorus is much less reactive than white phosphorus. It does not glow in the dark. Ignition temperature is 260ºC. It is polymeric, consisting of chains of P4 tetrahedra linked together in the manner as shown in Fig.

Black phosphorus has two forms a– black phosphorus and b – black phosphorus , a – Black phosphorus is formed when red phosphorus is heated in a sealed tube at 803 K. It can be sublimed in air and has opaque monoclinic or rhombohedral crystral. It does not oxidise in air.b – Black phosphorus is prepared by heating white phosphorus at 473 K under high pressure. It does not burn in air upto 673 K.

b - black phosphorus is a good conductor of electricity whereas -Black phosphorus is non-conductor.

b - black phosphorus has layered structure like graphite. The distance between the two layers is found to be 3.68 Å.

Density : White phosphorus= 1.83 ; Red phosphorus = 2.20 ; Black phosphorus  = 2.70 gm/cc ; 

As polymerisation increases compactness increases and therefore, density increases.

Reactivity of the various allotropic forms of phosphorus towards other substances decreases in the order:

white > red > black, the last one being almost inert i.e. most stable.

Apart from their reactivity difference, all the forms are chemically similar.

7. Compounds of Phosphorous


Phosphine :

Preparation :

(i) Phosphine is prepared by the reaction of calcium phosphide with water or dilute HCl.

 Ca3P2 + 6 H2®  3Ca(OH)2 + 2 PH3

Ca3P2 + 6HCl  ®  3CaCl2 + 2PH3

(ii) In the laboratory, it is prepared by heating white phosphorus with concentrated NaOH solution in an inert atmosphere of CO2.

P4 + 3 NaOH + 3 H2®  PH3 + 3 NaH2PO2   

                                                              (sodium hypophosphite)

Pure PH3 is non inflammable but becomes inflammable owing to the presence of P2H4 or P4 vapours. For removal of impurify, it is absorbed in HI to form phosphonium iodide (PH4I) which on treating with KOH gives off phosphine.

PH4I + KOH ®  KI + H2O + PH3

Properties :

(i) It is a colourless gas with a slightly garlic or rotten fish smell and is highly poisonous. It explodes in contact with traces of oxidising agents like HNO3, Cl2 and Br2  vapours.

(ii) It is slightly soluble in water but soluble in CS2 and other organic solvents. The solution of PH3 in water decomposes in presence of light giving red phosphorus and H2.

(iii)When absorbed in copper sulphate or mercuric chloride, the corresponding phosphides are obtained.

3CuSO4 + 2PH3 ® Cu3P2 ¯ (black) + 3H2SO4

3HgCl2 + 2 PH3 ® Hg3P2 ¯ (brownish black) + 6 HCl

(iv) Phosphine on heating at 150ºC burns forming H3PO4 

PH3 + 2O2 ® H3PO4           

(v) Phosphine is weakly basic and like ammonia, gives phosphonium compounds with acids e.g.,

 PH3 + HBr ® PH4Br

Phosphonium compounds are obtained when anhydrous phosphine reacts with anhydrous halogen acids (not in aqueous solution).

Uses :

The spontaneous combustion of phosphine is made to use in Holme’s signals. Containers containing calcium carbide and calcium phosphide are pierced and thrown in the sea when the gases evolved burn and serve as a signal.

It is also used in the production of smoke screens. Calcium phosphide reacts with water producing phosphine  which burns in air to give clouds of phosphorus pentaoxide and that acts as smoke screens. 

8. Phosphorus halides

Phosphorus Halides

Phsophorus forms two types of halides , PX3 [X = F , Cl , Br, I] and PX5 [X = F , Cl , Br]

(a) Phosphorus Trichloride :

Preparation :

(i)  .It is obtained by passing dry chlorine over heated white phosphorus

 P4 + 6 Cl2 ®  4 PCl3  

(ii) It is also obtained by the action of thionyl chloride with white phosphorus.

 P4 + 8 SOCl2  ®  4 PCl3 + 4 SO2  + 2 S2Cl2

Properties :

(i) It is a colourless oily liquid and hydrolyses in the presence of moisture.

 PCl3 + 3 H2 ® H3PO3 + 3 HCl

(ii) It reacts with organic compounds containing – OH group such as CH3COOH , C2H5OH.

  3 CH3COOH + PCl3  ® 3 CH3COCl + H3PO3 

  3 C2H5OH + PCl3  ® 3 C2H5Cl +H3PO3 

(b) Phosphorus pentachloride :

Preparation :

Phosphorus pentachloride is prepared by the reaction of white phosphorus with excess of dry chlorine.

P4 + 10 Cl2  ® 4 PCl5 

It can also be prepared by the action of SO2Cl2 on phosphorus.

P4 + 10 SO2Cl2  ® 4 PCl5 + 10 SO2

 Properties :

(i) PCl5 is a yellowish white powder and in moist air , it hydrolyses to POCl3 and finally gets converted to phosphoric acid.

 PCl5 + H2®  POCl3 + 2 HCl

 POCl3 + 3 H2® H3PO4 + 3 HCl

(ii) When heated it sublimes but decomposes on stronger heating.

PCl5  https://www.edumple.com/media/Images/CkEditor/%20Tarun%20Jaiswal%20__2379/598099.png  PCl3 + Cl2  

(iii) It reacts with organic compounds containing – OH group converting them to chloro derivatives.

C2H5OH + PCl5  ® C2H5Cl + POCl3 + HCl

CH3COOH + PCl5  ® CH3COCl + POCl3 + HCl

(iv) PCl5 on heating with finely divided metals give corresponding chlorides.

2 Ag + PCl5  ® 2 AgCl + PCl3

Sn + 2 PCl5   ®  SnCl4 + 2 PCl3

It is used in the synthesis of some organic compounds , e.g., C2H5Cl , CH3COCl.

9. Oxoacids of phosphorus

Oxides of phosphorus

(a) Phosphorus trioxide (P2O3) :

It is dimeric and has formula P4 O6

Preparation :

It is prepared by burning phosphorus in a limited supply of oxygen.

P4 + 3O(limited supply of oxygen)  P4O6

Properties :

(i) It is colourless crystalline solid having melting point 23.8oC and boiling point 178oC.

(ii) It dissolves in cold water to form phosphorus acid and in hot water liberating PH3.

 P4O6 + 6H2O (cold) ® 4H3PO3 

P4O6 + 6H2O (hot) ® 3H3PO4 + PH3

(iii) It burns in Cl2 gas forming phosphorus oxytrichloride (POCl3) and phosphoryl chloride (PO2Cl)

P4O6 + 4Cl2 ® 2POCl3 + 2PO2Cl

(b) Phosphorus pentaoxide (P2O5) :

It is dimeric and has the formula P4O10.

Preparation :

It is obtained by burning phosphorus in excess air.

P+ 5O2 ® P4O10


(i) It is a white powder ,acidic in nature and is the anhydride of orthophosphoric acid.

(ii)  It sublimes on heating at 250oC.

(iii)  It dissolves in water with hissing sound forming metaphosphoric acid and finally orthophosphoric acid.

P4O10 + 2H2® 4HPO3

4HPO3 + 2H2® 2H4P2O7

2H4P2O7 + 2H2® 4H3PO4

(iv) It dehydrates concentrated H2SO4 and concentrated HNO3 to SO3 and N2O5 respectively.

4HNO3 + P4O10 https://www.edumple.com/media/Images/CkEditor/%20Tarun%20Jaiswal%20__2379/532295.png4HPO3 + 2N2O5

2H2SO4 + P4O10 https://www.edumple.com/media/Images/CkEditor/%20Tarun%20Jaiswal%20__2379/559504.png 4HPO3 + 2SO3


1. For drying acidic gases.  

2. As a dehydrating agent

3. For the preparation of SO3 and N2O5.

4. For the preparation of phosphoric acid.

Oxoacids of Phosphorus :

Phosphorus forms a number of a oxoacids as given in following Table :


The structures of some of oxo-acids are as given below :



In oxoacids phosphorus is tetrahedrally surrounded by other atoms. All these acids contain one P=O and at least one P–OH bond. The oxoacids in which phosphorus has lower oxidation state (less than +5) contain in addition to P=O and P–OH bonds ,either P–P (e.g., in H4P2O6) or P–H (e.g., in H3PO2) bonds but not both. These acids in + 3 oxidation state of phosphorus tend to disproportionate to higher and lower oxidation states. For example, orthophosphorus acid (or phosphorus acid) on heating disproportionates to give orthophosphoric acid (or phosphoric acid) and phosphine.

 4 H3PO3  ® 3 H3PO4  + PH3

The acids which contanin P – H bond have strong reducing properties. Thus , hypophorous acid is a good reducing agent as it contains two P – H bonds and reduces , for example, AgNO3 to metallic silver.

4 AgNO3 + 2 H2O + H3PO2 ®4 Ag + 4 HNO3 + H3PO4 

These P–H bonds are not ionisable to give H+ and do not play any role in basicity. Only those H atoms which are attached with oxygen in P–OH form are ionisable and cause the basicity. Thus , H3PO3 and H3PO4 are dibasic and tribasic, respectively as the structure of H3PO3has two P – OH bonds and H3PO4 three.

10. Group 16 elements

Group 16 Elements : The Oxygen family

Oxygen, sulphur, selenium, tellurium and polonium constitute group 16 of the periodic table. This is sometimes known as group of chalcogens the ore forming elements because a large number of metals ores are oxides or sulphides.

Electronic Configuration : The elements of group 16 have six electrons in the outermost shell and have ns2 np4 general valence shell electronic configuration.

Atomic and Ionic Radii : Due to increase in the number of shells , atomic and ionic radii increase from top to bottom in the group. The size of oxygen atoms is however, exceptionally small .

Ionisation Enthalpy : Ionisation enthalpy decreases down the group. It is due to increase in size. However, the element of this group have lower ionisation enthalpy values compared to those of group 15 in the corresponding periods. This is due to the fact that group 15 elements have extra stable half-filled p orbitals electronic configurations.

Electron Gain Enthalpy : Because of the compact nature of oxygen atom, it has less negative electron gain enthalpy than sulphur. However from sulphur onwards the value again becomes less negative upto polonium.

Electronegativity : Next to fluorine, oxygen has the highest electronegativity value amongst the elements. Within the group, electronegativity decrease with an increase in atomic number. This indicates that the metallic character increases from oxygen to polonium.

Physical Properties : Oxygen and sulphur are non-metal, selenium and tellurium metalloids, whereas polonium is a metal. Polonium is radioactive and is short lived (Half-life 13.8 days). The melting and boiling points increase with an increase in atomic number down the group. The larger difference between the melting and boiling points of oxygen and sulphur may be explained on the basis of their atomicity; oxygen exist as diatomic molecules (O2) whereas sulphur exists as polyatomic molecule (S8).

Catenation : Tendency for catenation decreases down the group. This property is prominently displayed by sulphur (S8). The S—S bond is important in biological system and is found in some proteins and enzymes such as cysteine.

Selenium has unique property of photo conductivity and is used in photocopying machines and also a decolouriser of glass.

Atomic & Physical Properties

Chemical Properties :   

Oxidation states and trends in chemical reactivity :

The elements of group 16 exhibit a number of oxidation states. The stability of -2 oxidation state decreases down the group. Polonium hardly shows -2 oxidation states. Since electronegativity of oxygen is very high, it shows only negative oxidation states as -2 except in the case of OF2 where its oxidation states is + 2. Other elements of the group exhibit + 2 + 4 + 6 oxidation states but + 4 and + 6 are more common. Sulphur, selenium and tellurium usually show + 4 oxidation in their compounds with oxygen and +6 oxidations state with fluorine. The stability of +6 oxidation state decreases  down the group and stability of + 4 oxidation state increases (inert pair effect). Bonding in + 4 and + 6 oxidation states are primarily covalent.

HNO3 oxidises sulphur to H2SO4 (S + VI) but only oxidises selenium to H2SeO3 (Se + IV) as the atoms are smaller and there is poor shielding of 3d electrons as a result the electrons are held more tightly with nucleus.

Anomalous behaviour of oxygen :

The anomalous behaviour of oxygen, like other member of p-block  present in second period is due to its small size and high electronegativity. One typical example of effects of small size and high electronegativity is the presence of strong hydrogen  bonding in H2O which is not found in H2S.

The absence of d orbitals in oxygen restricts its covalency to four and in practice, rarely increases beyond two. On the other hand, in case of other elements of the group, the valence shell can be expanded and covalence exceeds four.

(i) Reactivity with hydrogen : All the elements of group 16 form hydrides of the type H2E (E = S, Se, Te, Po). Some properties of hydrides are given in Table. Their acidic character increases from H2O to H2Te. The increase  in acidic character can be understood in terms of decrease in bond (H-E) dissociation enthalpy down the group. Owing to the decrease in bond (H-E) dissociation enthalpy down the group , the thermal stability of hydrides also decreases from H2O to H2Po. All the hydrides except water possess reducing property and this property increases from H2S to H2Te.

Table : Properties of Hydrides of Group 16 Elements

(ii) Reactivity with oxygen : All these elements form oxides of the EO2 and EO3 types where E = S, Se, Te or Po. Ozone (O3) and sulphur dioxide (SO2) are gases while selenium dioxide (SeO2) is solid. Reducing property of dioxide decreases from SO2 to TeO2 ; SO2 is reducing while TeO2 is an oxidising agent. Besides EO2 type sulphur, selenium and tellurium also form EO3 type oxides (SO3, SeO3, TeO3). Both types of oxides are acidic in nature.

(iii) Reactivity toward the  halogens : Elements of group 16 form a larger number of halides of the type EX6, EX4 and EX2  where E is an element of the group  and X is an halogen. The stabilities of the halides decrease in  the order F > Cl > Br > l. Amongst hexahalides, hexafluorides are the only stable halides. All hexafluorides are gaseous in  nature. They have octahedral structure. Sulphur hexafluoride SF6 is exceptionally  stable for steric reasons.

Amongst tetrafluorides, SF4 is a gas , SeF4 liquid  and TeF4 a solid.

All elements except selenium form dichlorides and dibromides.The well known monohalides are dimeric in nature, Examples are S2F2, S2Cl2, S2Br2, Se2Cl2 and Se2Br2. These dimeric halides undergo disproportionation as given below :

2Se2Cl2  ® SeCl4 + 3Se.

11. Dioxygen


It differs from the remaining elements of the VIth group because  of the following properties.

(A) small size

(B) high electronegativity and

(C) non-availability of d-orbitals.

Preparation :

(i) By thermal decomposition of oxides of metals.

2 HgO  https://www.edumple.com/media/Images/CkEditor/%20Tarun%20Jaiswal%20__2379/261239.png 2 Hg + O2

2 Ag2O  https://www.edumple.com/media/Images/CkEditor/%20Tarun%20Jaiswal%20__2379/253428.png 4 Ag + O2 

(ii) By thermal decomposition of oxygen rich compounds.                                                 

KClO3   https://www.edumple.com/media/Images/CkEditor/%20Tarun%20Jaiswal%20__2379/323190.png  2 KCl + 3O2 (laboratory method)

(iii) 2H2O2(aq.) https://www.edumple.com/media/Images/CkEditor/%20Tarun%20Jaiswal%20__2379/169799.png 2H2O(l) + O2(g)

(iv) Industrial method :

(a) Electrolysis of water leads to the release of hydrogen at the cathode and oxygen at the anode.

 (b) Oxygen is obtained by liquification of air and then its fractional distillation.

Physical properties :

Colourless , odourless and tasteless gas. It is paramagnetic and exhibits allotropy. Three isotopes of oxygen are https://www.edumple.com/media/Images/CkEditor/%20Tarun%20Jaiswal%20__2379/19163.pnghttps://www.edumple.com/media/Images/CkEditor/%20Tarun%20Jaiswal%20__2379/31809.png and https://www.edumple.com/media/Images/CkEditor/%20Tarun%20Jaiswal%20__2379/934326.png. Oxygen does not burn but is a strong supporter of combustion.

Chemical properties :

(i) Reaction with metals :

2Ca + O2  ® 2CaO  

4Al + 3O2 ® 2Al2O3

(ii) Reaction with non-metals :

P4 + 5O®  P4O10

C + O2 ® CO2

 (iii) Reaction with compounds :

2ZnS + 3O®  2ZnO + 2SO2 

CH4 + 2O2 ®CO2 + 2H2O

2SO2 + O2 https://www.edumple.com/media/Images/CkEditor/%20Tarun%20Jaiswal%20__2379/965709.png 2SO3

4HCl + O2  https://www.edumple.com/media/Images/CkEditor/%20Tarun%20Jaiswal%20__2379/250142.png 2Cl2 + 2H2O

Note : It has been observed that its combination with other elements is often strongly exothermic which helps in sustaining the reaction. However, to initiate the reaction, some external heating is required as bond dissociation enthalpy of oxygen-oxygen double bond is high (493.4 kJ mol–1).

Use : 

1. Oxygen mixed with helium or CO2 is used for artificial respiration.

2. Liquid oxygen (with combustion fuel hydrazine) is used as oxidising agent in rocket fuels.

3. Oxygen is used for production of oxy-hydrogen or oxy-acetylene flames employed for cutting and welding.

4. Pure dioxygen is used to convert pig iron into steel in the basic oxygen process which are kaldo and LD process.

12. Oxides


(i)  Acidic oxides :

The covalent oxides of non-metal are usually acidic; dissolves in water to produce solutions of acids e.g., CO2, SO2 , SO3, N2O5 , N2O3 , P4O6 , P4O10, Cl2O7, CrO3 ,  Mn2O7 etc. They are termed as acid anhydride.

Cl2O7 + H2O  ®  2 HClO4

Mn2O7 + H2O  ® 2 HMnO4

(ii) Basic oxides :

Metallic oxides are generally basic oxides. They either dissolve in water to form alkalies or combine with acids to form salts and water or combine with acidic oxides to form salts; e.g., Na2O, CaO. CuO, FeO, BaO etc.

Na2O + H2O  ® 2 NaOH

CaO + H2 ® Ca(OH)2

CuO + H2SO4   ® CuSO4 + H2O

The metallic oxides with lowest oxidation sate is the most ionic and the most basic but with increasing oxidation sate the acidic character increases e.g., CrO is basic, Cr2O3 amphoteric and CrO3 acidic.

(iii) Amphoteric Oxides :

 Many metals yield oxides which combine with both strong acid as well as strong bases e.g., ZnO, Al2O3, BeO, Sb2O3, Cr2O3, PbO, SnO, SnO2, Ga2O3 etc.

Cr2O3 + 2 NaOH ® Na2Cr2O4 + H2O

Cr2O3 + 3 H2SO4 ® Cr2(SO4)3 + 3 H2O

13. Ozone


O3 is an allotropic form of oxygen. At a height of about 20 Kms it is formed from atmoshperic oxygen in the presence of sunlight. This O3 layer protects the earth’s surface from an excessive concentration of ultra violet radiations.

Preparation :

It is prepared by passing silent electric discharge through a slow stream of pure and dry oxygen to prevent its decomposition.


DHV (298 K) = + 142 kJ mol–1

The product is known as ozonised oxygen. If concentration of O3 greater than 10% are required , a battery of ozonisers can be used , and pure ozone (bp 385 K) can be condensed in a vessel surrounded by liquid oxygen.

Properties :

Physical properties :

(1) It is a pale blue gas which forms a blue liquid and one solidification forms violet black crystals.

(2)  It has a strong fish – like smell

(3) It is slightly soluble in water but more in turpentine oil or glacial acetic acid or CCl4 .

(4) O3 molecule is diamagnetic but O3 ion is paramagnetic (1 unpaired e)

(5) It is explosive and unstable with respect to O2 as its decomposition into O2 results in the liberation of heats  and an increase in entropy.

Chemical Properties :

(1)  As Oxidising agent : Due to the ease with which it liberates atoms of nascent oxygen

(O3  ® O2 + O ), it acts as a powerful oxidising agent.

In acidic medium :

O3 + 2 H+ + 2e ® O2 + 2 H2O   SRP = + 2.07 V.

In alkaline medium :

O3 + H2O + 2e ® O2 + 2 OH   SRP = + 1..24 V

 Therefore , Ozone is a stronger oxidising agent in acidic medium.

With excess of potassium iodide solution buffered with a borate buffer, ozone liberates iodine which can be titrated against a standard solution of sodium thiosulphate. This is a quantitative method for estimating O3 gas.

It oxidises PbS to PbSO4 , MnO42– to MnO4 (basic medium) and [Fe(CN)6]4– to [Fe(CN)6]3–  (basic medium).

Note : With experimental facts it has been shown that nitrogen oxides (particularly nitric oxide) combine very rapidly with ozone and there is, thus, the possibility that nitrogen oxides emitted from the exhaust systems of supersonic jet aeroplanes might be slowly depleting the concentration of the ozone layer in the upper atmosphere.

NO(g) + O3 (g) ® NO2 (g) + O2 (g)

O–O bond length decreases in order : H2O2 (1.48 Å) > O3 (1.28 Å) > O2F2 (1.22 Å) > O2 (1.21 Å)

Uses :          

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

14. Sulphur - Allotropic forms


Allotropic Forms Of Sulphur :

Sulphur forms numerous allotropes of which the yellow rhombic  (a - sulphur) and monoclinic (b - sulphur) forms are the most important. The stable forms at room temperature is rhombic sulphur, which transforms to monoclinic sulphur when heated above 369 K.

Rhombic sulphur  (a - sulphur) :

This allotrope is Syellow in colour , melting point 385.8 K and specific gravity 2.06. Rhombic sulphur crystals are formed on evaporating the  solution of roll sulphur in CS2. It is insoluble in water but dissolved  to some extent in benzene, alcohol and ether. It is readily soluble in CS2 .

Monoclinic sulphur (b - sulphur) :

Its melting point is 393 K and specific gravity 1.98. It is soluble in CS2. This form of sulphur is prepared by melting rhombic sulphur in a dish and cooling till crust is formed. Two holes are made in the crust and the remaining liquid poured out. On removing the crust, colourless needle shaped crystals of b - sulphur are formed. It is stable above 369 K and transforms into a - sulphur below it . Conversely, - sulphur is stable below 369 K and transforms into b - sulphur above this.  At 369 K both the forms are stable. This temperature is called transition temperature.

Both rhombic and  monoclinic sulphur have S8 molecules these S8 molecules are packed to give different crystal structures. The S8 ring in both the forms is puckered and has a crown shape. The molecular dimensions are given in figure.

Several other modifications of sulphur containing 6-20 sulphur atoms per ring have been synthesised in the last two decades. In cyclo- S6, the ring adopts the chair form and the molecular dimension are as shown in fig. (b).

Sulphur melts to form a mobile liquid. As the temperature is raised the colour darkens. At 160ºC C8 rings break, and the diradicals so formed polymerize, forming long chains of up to a million atoms. The viscosity increases sharply, and continues to rise up to 200ºC. At higher temperatures chains break, and shorter chains and rings are formed, which makes the viscosity decrease upto 444ºC, the boiling point. The vapour at 200ºC consists mostly of S8 rings, but contains 1-2% of S2 molecules. At elevated temperature (~1000 K), S2 is the dominant species and is paramagnetic like O2, and presumably has similar bonding. S2 gas is stable upto 2200ºC.

15. Sulphur dioxide


Perparation :

(i)  Sulphur dioxide is formed together with a little (6-8%) sulphur trioxide when sulphur is burnt in air or oxygen:

S(s) + O2(g)  ®  SO2(g)

(ii)   In the laboratory it is readily generated by treating a sulphite with dilute sulphuric acid.

SO32 – (aq) + 2H+ (aq) ® H2O(l) + SO2(g)

(iii)  Industrially it is produced as a by- product of the roasting of sulphide ores.

4FeS2(s) + 11O2 (g) ®2Fe2O3(s) + 8SO2(g)

The gas after drying is liquefied under pressure and stored in steel cylinders.

Properties :

(i) Sulphur dioxide is a colorless gas with pungent smell and is highly soluble in water. It liquefies at room temperature under a pressure of two atomsphere and boils at 263 K.

(ii) Sulphur dioxide, when passed through water, forms a solution of sulphurous acid.

 SO2(g) + H2O(l)  H2 SO3 (aq)

(iii)It reacts readily with sodium hydroxide solution, forming sodium sulphite which then reacts with more sulphur dioxide to form sodium hydrogen sulphite.

2NaOH + SO2 — Na2SO3 + H2O

  Na2SO3 + H2O + SO2 ® 2NaHSO3 

(iv) In it reaction with water and alkalies, the behaivour of sulphur dioxide is very similar to that of carbon dioxide.

Sulphur dioxide reacts with chlorine in the presence of charcoal (which acts as a catalyst) to give sulphuryl chloride, SO2Cl2 It is oxidised to sulphur trioxide by oxygen in the presence of vanadium (v) oxide catalyst.

SO2(g) + Cl2(g) ® SO2Cl2(l)

2SO2 (g) + O2 (g)  2SO3(g)     Addition reactions

(v) When moist, sulphur dioxide behaves as a reducing agent. For example it converts iron (III) ions to irons (II) ions and decolourises acidified potasssium permanganate (VII) solution; the latter reaction is a convenient test for the gas.

2Fe3+ + SO2 + 2H2® 2Fe2+ + SO42–  + 4H+ 

5SO2 + 2MnO4 + 2H2O ®  5SO42–  + 4H+   + 2Mn2+

(vi)  Bleaching action : 

SO2 + 2H2® H2SO4 + 2H

 coloured matter  colourless matter

Bleaching is through reduction but it is temporary.

It is a more powerful reducing agent in alkaline medium than in acidic medium.

Uses :          

Sulphur dioxide is used (i) in refining petroleum and sugar (ii) in bleaching wool and silk and (iii) as an anti- chlor, disinfecatant and preservation. Sulphuric acid, sodium hydrogen sulphite and calcium hydrogen sulphite (industrial chemicals) are manufactured from sulphur dioxide. Liquid SO2 is used as a solvent to dissolve a number of organic and inorganic chemicals.

16. Oxoacids of sulphur


Sulphur forms a number of oxoacid such as H2SO3, H2S2O4, H2S2O5, H2S2O6 (x = 2 to 5,) H2SO4, H2S2O7,  H2SO8.Structures of some important oxoacids are shown in figure.

17. Sulphuric acid


Manufacture :

Sulphuric acid is manufactured by the contact process which involves three steps :

(i) burning of sulphur or sulphide ores in air to generate SO2

(ii) conversion of SO2 to SO3 by the reaction with oxygen in the presence of a catalyst (V2O5), and

(iii) absorption of SO2 in H2SO4 to give Oleum (H2S2O7)

The SO2 produced is profiled by removing dust and other impurities such as arsenic compounds.

The key step in the manufacture of H2SO4 is the catalytic oxidation of SO2 with O2 to give SO3 in the presence of V2O5 (catalyst).

2SO2(g) +  O2(g)  2SO3(g) DrH = – 196.6 kJ mol–1.

The reaction is exothermic reversible and the forward reaction leads to a decrease in volume. Therefore, low temperature and high pressure are the favourable conditions for maximum yield. But the temperature should not be very low other wise rate of reaction will become slow.

In practice the plant is operated at a pressure of 2 bar and a temperature of 720 K. The SO3 gas from the catalytic converter is absorbed in concentrated H2SO4 to produce oleum. Dilution of oleum with water gives H2SO4 of the desired concentration. In the industry two steps are carried out simultaneously to make the process a continuous one and also to reduce the cost.

SO3 +  H2SO4  ®    H2S2O7 


The sulphuric acid obtained by Contact process is 96-98% pure.

Properties :

Sulphuric acid is a colourless, dense, oily liquid with a specific gravity of 1.84 at 298 K. The acid freezes at 283 K and boils at 611 K. It dissolves in water with the evolution of a larger quantity of heat. The chemical reaction of sulphuric acid are as a result of the following characteristics : (a) low volatility (b) strong acidic character (c) strong affinity for water and (d) ability to act as an oxidising agent in aqueous solution,

(i) Sulphuric acid ionises in two steps.

H2SO4(aq)  + H2O(l) ®  H3O+ (aq) + HSO4 (aq)

Ka1 = very larger (Ka1 > 10)

H2SO4 (aq)  + H2O(l) ® H3O+ (aq) + SO42– (aq)

Ka2 = 1.2 × 10–2 

The larger value of Ka1 (Ka1 > 10) means that H2SO4 is largely dissociated into H+ and HSO4. Greater the value of dissociation constant (Ka) the stronger is the acid.

(a) The acid forms two series of salts : normal sulphates (such as sodium sulphate and copper sulphate   and acid sulphate (e.g., sodium hydrogen sulphate)

(b) Decomposes carbonates and bicarbonates in to C.

Na2CO3 +  H2SO4  ®  Na2SO4 +  H2O +  CO2

NaHCO3 + H2SO4 ®  NaHSO4 +  H2O + CO2

(c) Sulphuric acid, because of its low volatility can be used to manufacture more volatile acid from their corresponding salts.

2MX +  H2SO4 ® 2HX +  M2SO4 (X = F, Cl, NO3

NaCl + H2SO4 ®  NaHSO4 + HCl

                                 (M = Metal)

(ii) Concentrated sulphuric acid is a strong dehydrating agent. Many wet gases can be dried by passing them through sulphuric acid, provided the gases do not react with the acid. Sulphuric acid removes water from organic compound; it is evident by its charring action on carbohydrates.

C12H22O11  12C + 11H2O

H2C2O4CO + CO2

(iii) Hot concentrated sulphuric acid is moderately strong oxidising agent. In this respect it is intermediate between phosphoric and nitric acids. Both metals and non-metals are oxidised by concentrated sulphuric acid, which is reduced to SO2.

Cu + 2H2SO4 (concentrated)  ® CuSO4 + 2H2O

3S  + 2H2SO4 (concentrated)  ® 3SO2 + 2H2O

C  + 2H2SO4 (concentrated)  ® CO2 + 2SO2 + 2H2O

Uses :

Sulphuric acid is a very important industrial chemical. A nation’s industrial strength can be judged by the quantity of sulphuric acid it produces and consumes .It is needed for the manufacture of hundreds of other compounds also in many industrial processes .The bulk of sulphuric acid produced is used in the manufacture of fertilisers (e.g., ammonium sulphate, superphosphate). Other uses are in : (i) petroleum refining (ii) manufacture of pigment, paints and dyestuff intermediates (iii) detergent industry (iv) metallurgical applications (e.g., cleansing metal before enameling, electroplating and galvanising) (v) storage batteries (vi) in the manufacture of nitrocellulose products and (vii) as a laboratory reagent.

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