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Industrial Chemistry

Industrial Chemistry Synopsis

Synopsis


Some important nonmetals:
Bonding in Carbon

  • Carbon atom has four electrons in its outermost shell.

     
  • It requires four electrons to achieve the stable, 8 electron, inert gas electron arrangement.
  • Carbon atoms can achieve the inert gas electron arrangement only by sharing of the electrons. Hence, carbon always forms covalent bonds.
  • The valency of carbon is four since one carbon requires 4 electrons to achieve the nearest inert gas configuration. Thus, we can say that carbon is tetravalent.
  • The four valencies of carbon are usually represented by putting four short lines around the symbol of carbon, C.

      

Allotropes of Carbon

  • The various physical forms in which an element can exist are called the allotropes of the element.
  • Three allotropes of carbon are:

     
  • Diamond
    • In diamond, each carbon atom is bonded to four other carbon atoms, forming a three dimensional structure.
    • The rigid structure of diamond makes it a very hard substance.
    • It is a non-conductor of electricity since there are no free electrons in a diamond crystal.
    • It can be synthesised by subjecting pure carbon to very high pressure and temperature.
  • Graphite
    • In graphite, each carbon atom is bonded to three other carbon atoms in the same plane, giving a hexagonal array.
    • One of the bonds is a double bond and thus the valency of the carbon is satisfied.
    • Graphite structure is formed by the hexagonal arrays being placed in layers, one above another.
    • Graphite is smooth and slippery.
    • It is a very good conductor of electricity due to the presence of free electrons. 

  • Fullerene:
    • It is an allotrope of carbon containing clusters of 60 carbon atoms joined together to form spherical molecules.
    • There are 60 carbon atoms in a molecule of buckminsterfullerene, so its formula is C60.
    • The allotrope was named buckminsterfullerene after the American architect Buckminster Fuller.

                          

 

Versatile Nature of Carbon

  • The number of carbon compounds already known at present is more than 5 million.
  • Every day more new compounds are isolated or prepared in the laboratories.
  • The two characteristic properties of the carbon element which leads to the formation of a very large number of organic compounds are
  1. Catenation: The property of carbon element due to which its atoms can join one another to form long carbon chains is called catenation.
    A)  Straight chain of carbon atoms



    B)  Branched chain of carbon atoms


          C)  Closed chain or ring chain of carbon atoms


            

      

ii. Tetravalency: Carbon has a valency of four. So, it is capable of bonding with four other atoms of carbon or atoms of some other mono-valent element. Compounds of carbon are formed with oxygen, nitrogen, hydrogen, sulphur, chlorine and many other elements, giving rise to compounds with specific properties which depend on the elements other than the carbon present in the molecule.

Compounds of Nitrogen:
Ammonia
Molecular formula: NH3
Relative molecular mass: 17 am

Lewis diagram or dot diagram

 Occurrence

  • Free State:
    Ammonia is present in small amounts in air and in traces in natural water.
    Combined state:
    Ammonia occurs in the combined form in compounds such as ammonium chloride and ammonium sulphate.
 Forms of Ammonia

              i.   Gaseous ammonia
             ii.   Liquid ammonia
            iii.    Liquor ammonia fortis
            iv.    Laboratory bench reagent

Preparation of Ammonia Gas
 Laboratory Preparation

  1. From Ammonium Chloride
    Reactants: Ammonium chloride (NH4Cl) and calcium hydroxide [Ca(OH)2] in the ratio of 2:3 by weight.
    Reaction:
    2NH4Cl + Ca(OH)2  → CaCl2 + 2H2O + 2NH3

    Diagram:

     

 

Drying of Ammonia Gas

  1.  The gas is passed through a drying tower containing lumps of quicklime (CaO).
  2.  Drying agents such as conc. sulphuric acid, phosphorous pentoxide and calcium chloride are not used because ammonia, being basic, reacts with them.
    2NH3 + H2SO4 → (NH4)2SO4
    6NH3 + P2O5 + 3H2O → 2(NH4)3PO4
    CaCl2 + 4NH3 → CaCl2.4NH3

Collection
Ammonia gas is collected by the downward displacement of air because it is

  1.  Lighter than air (VD of NH3 = 8.5 and that of air = 14.4)
  2.  Highly soluble in water (so, it cannot be collected over water)

2. From Metal Nitrides
It can also be prepared by the action of warm water on nitrides of metals such as magnesium or aluminium.
Reaction:
AlN + 3H2O  →  Al(OH)3 + NH3

                   Or                             

Mg3N2 + 6H2O  →  3Mg(OH)2 + 2NH3

Manufacture of Ammonia (Haber’s Process)
Reactants:
Nitrogen and hydrogen in the ratio of 1:3 by volume.

Reaction:
N2 + 3H2  ⇌ 2NH3 + Heat

Sources of Reactants
Nitrogen: It is obtained by fractional distillation of liquid air.
Hydrogen: It is obtained from water gas (Bosch process) or natural gas.
Favourable Conditions
Temperature: Optimum temperature is 450−500°C.
Pressure: Above 200 atm.
Catalyst: Finely divided iron
Promoter: Traces of molybdenum or Al2O3

Chemical Properties of Ammonia

 

Uses of Ammonia

  1.  Liquid ammonia is used as a refrigerant in ice plants.
  2.  Aqueous NH3 dissolves fat and grease, so it is used
    a. To remove grease and sweat stains from clothes
    b. For cleaning tiles, windows etc.
  3. Used as a laboratory reagent in qualitative analysis because it produces characteristic coloured metallic hydroxide precipitates.
  4. Ammonia is used in the manufacture of
    1. Nitrogenous fertilisers such as ammonium sulphate and diammonium hydrogen phosphate
    2. Explosives such as ammonium nitrate
    3. Sodium carbonate by the Solvay process
    4. Nitric acid by the Ostwald process

Nitric Acid
Molecular formula: HNO3
Relative molecular mass: 63 amu

  

  • Nitric acid was initially called ‘aqua fortis’.
  • In 1658, Glauber obtained nitric acid by distilling nitre with sulphuric acid.

Occurrence
Free state: Nitric acid is found in rainwater in traces after lightning.
Combined state: As its salts in minerals
Examples: Chile saltpetre [NaNO3], Bengal saltpetre [KNO3] or nitre

Laboratory Preparation of Nitric Acid
Reactants:
Obtained by distilling conc. sulphuric acid with potassium KNO3 (nitre) or sodium NaNO3.

Reactions:
begin mathsize 12px style KNO subscript 3 space plus space straight H subscript 2 SO subscript 4 space rightwards arrow with less than 200 to the power of ring operator straight c on top space space space KHSO subscript 4 space plus space HNO subscript 3
NaNO subscript 3 space plus straight H subscript 2 SO subscript 4 space rightwards arrow with less than 200 to the power of ring operator straight c on top NaHSO subscript 4 space plus space HNO subscript 3 end style

Collection

  • Vapours of nitric acid are condensed into a light yellow liquid by chilling the receiver with running cold water.
  • Pure nitric acid is colourless, while the nitric acid obtained in the laboratory is slightly yellowish-brown due to its decomposition, resulting in the formation of reddish-brown nitrogen dioxide.
  • The yellow brown tinge in the acid can be removed by
    1. Bubbling of carbon dioxide, which oxidises nitrogen dioxide to nitric acid
    2. Dilution with water, which causes dissolution of water-soluble nitrogen dioxide

Precautions

  • Glass apparatus is used because nitric acid vapours corrode rubber and cork.
  • Conc. HCl is not used because it is volatile, and hence, nitric acid vapours will carry HCl vapours.
  • The temperature of the reaction should not exceed 200°C because sodium sulphate formed at a higher temperature forms a hard crust which sticks to the walls of the retort and is difficult to remove.

            begin mathsize 12px style NaNO subscript 3 space plus space NaHSO subscript 4 space end subscript rightwards arrow with less than 200 to the power of ring operator straight c on top Na subscript 2 SO subscript 4 space plus space HNO subscript 3 end style

Manufacture of Nitric Acid
In 1914, a German chemist Ostwald developed the Ostwald process to manufacture nitric acid.

Step 1: Catalytic oxidation of ammonia
A mixture of dry air and dry ammonia in the ratio of 10:1 by volume is compressed and then passed into a platinum gauze which acts as a catalyst at about 800°C.
begin mathsize 10px style 4 NH subscript 3 space end subscript plus space 5 straight O subscript 2 space space rightwards arrow with Platinum comma 800 to the power of ring operator straight c on top 4 NO space plus space 6 straight H subscript 2 straight O space plus space Heat end style

Step 2: Oxidation of nitric oxide
Nitric oxide combines with oxygen to form nitrogen dioxide at about 50°C.
begin mathsize 12px style 2 NO space plus straight O subscript 2 space rightwards arrow with 50 to the power of ring operator straight c on top 2 NO subscript 2 end style

Step 3: Absorption of nitrogen dioxide in water
Nitrogen dioxide and oxygen present in the air react with water to form nitric acid.
4NO2 + 2H2O + O2 → 4HNO3

Nitric acid obtained is concentrated above 50%. On further distillation, 68% nitric acid is produced.

Properties of Nitric Acid

  • Physical Properties
    Colour, odour and taste: Pure acid (98% conc.) is colourless, suffocating and sour to taste.
    Physiological nature: Non-poisonous, highly corrosive
    Density: Heavier than water, with specific gravity 1.54 g/cm3
    Solubility: Highly soluble in water
    Boiling point: 86°C
    Freezing point: −42°C
  • Chemical Properties

     

 Uses of Nitric Acid

  • To etch designs on copper and brassware because it acts as a solvent for several metals, except the noble metals
  • To purify gold from impurities of Cu, Ag and Zn which dissolve in nitric acid
  • As a rocket fuel oxidant
  • In preparation of fertilisers such as Ca(NO3)2 and NH4NO3
  • In preparation of aqua regia, which dissolves noble metals

Industrial Uses
In the manufacture of

  1.  Explosives such as TNT
  2.  Synthetic fertilisers such as artificial silk, nylon, plastics etc.
  3. Important compounds such as nitrates of potassium, ammonium, silver etc.
  4. Dyes, drugs and perfumes

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

White phosphorus

  • It is a translucent white waxy solid. It is poisonous; it is soluble in carbon disulphide and glows in dark (chemiluminescence). It dissolves in boiling NaOH solution in an inert atmosphere giving PH3, but it is insoluble in water.

             P4 + 3NaOH + 3H2O → PH3 + 3NaH2PO2
                                                         (sodium hypohoshite)                    

  • White phosphorus having less stability because of that it is more reactive than the other solid phases under normal conditions because of angular strain in the P4 molecule where the angles are only 60°. It readily catches fire in air to give dense white fumes of P4O10.
    P4 + 502 → P4O10 
  • It made up of individually separate tetrahedral P4molecule.

Red phosphorus

  • It is obtained by heating white phosphorus at 573K in an inert atmosphere for several days.
  • After the red phosphorus is heated under high pressure, a series of phases of black phosphorus are formed.
  • Red phosphorus possesses iron grey lustre.
  • It is odourless, nonpoisonous and insoluble in water as well as in carbon disulphide.
  • Red phosphorus is polymeric, consisting of chains of P4 tetrahedral linked together in the manner.

Black phosphorus

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

Preparation

  • The reaction of calcium phosphide with water or dilute HCl forms Phosphine is prepared.

           Ca3P2 + 6H2O → 3 Ca(OH)2 +2 PH3
           Ca3P2 + 6HCI → 3 CaCI2 +2 PH3

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

           P4 + 3NaOH + 3H2O → PH3 + 3NaH2PO2
                                                         (sodium hypohoshite)

  • It is non inflammable when pure, but becomes inflammable owing to the presence of P2H4 or P4 vapours.
  • To purify it from the impurities, it is absorbed in HI to form phosphonium iodide (PH4I) which on treating with KOH gives off phosphine.

           PH4I+ KOH → KI + H2O + PH3

Properties
Physical properties

  • It is a colourless gas with rotten fish smell and is highly poisonous.
  • It explodes in contact with traces of oxidising agents like HNO3, Cl2 and Br2 vapours.

Chemical properties

  • It is slightly soluble in water. The solution of PH3 in water decomposes in presence of light giving red phosphorus and H2.
  • When absorbed in copper sulphate or mercuric chloride solution, the corresponding phosphides are obtained.
    3CUSO4 + 2 PH3 → Cu3P2 +3H2SO4
    3HgCI2 + 2 PH3 → Hg3P2 +6HCI
  • Phosphine is weakly basic and like ammonia, gives phosphonium compounds with acids.
    e.g.,  PH3 + HBr → PH4Br

Uses

  • The spontaneous combustion of phosphine is technically used 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 smoke screens.

Sulphur and its compounds
Sulphur – allotropic forms
Sulphur forms variety of allotropes. The most common allotropes are yellow rhombic and monoclinic – sulphur. Rhombic sulphur is more stable at room temperature. It gets transforms to monoclinic sulphur when heated above 369 K.

Rhombic sulphur

  • This allotrope is yellow in colour, m.p. is around 385.8 K and specific gravity is 2.06.
  • Rhombic sulphur crystals are formed when the solution of roll sulphur in CS2 is evaporated.
  • It is insoluble in water but dissolves to some extent in benzene, alcohol and ether. It is more soluble in CS2.

Monoclinic sulphur

  • Its m.p. 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. After removing crust, colourless needle shaped crystals of sulphur are formed.
  • It is stable above 369 K and transforms into sulphur below 369 K.
  • Also we can say that the sulphur is stable below 369 K and transforms into sulphur above this. At 369 K both the forms are stable. This temperature is called transition temperature.
  • Rhombic as well as 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

  • In cyclo-S6 form, the molecule is in chair shape.