oxygen family - theory notes with illustrative examples [unlocked by www.freemypdf.com]

15
IIT-JEE Via Group Elements 1 VIA GROUP ELEMENTS Introduction Oxygen, sulphur, selenium, Tellurium and polonium are the elements of VIA group. They have ns 2 np 4 outer electron configuration. They belong to ‘p’ block in the long form of the periodic table. The first four elements are collectively known as Chalcogens (means mineral forming) since many metals occur as oxides or sulphides. Polonium is a radioactive element and short lived. It was first isolated by Madam Curie from an ore of Uranium called Pitchblende. Curie named this element as Polonium in honour of her native land, Poland. The elements of this group have six electrons in their outer shell. These electrons are described as valence electrons. In these six electrons two electrons are in ‘s’ orbital and four electrons are in ‘p’ orbitals. In these four electrons two are unpaired. These unpaired electrons generally take part in chemical bonding. All these elements require two electrons to attain a stable inert gas configuration. Electron Configuration of VIA Group Elements Element Symbol Atomic Number Electron Configuration Oxygen O 8 [He] 2s 2 sp 4 Sulphur S 16 [Ne] 3s 2 3p 4 Selenium Se 34 [Ar] 3d 10 4s 2 4p 4 Tellurium Te 52 [Kr] 4d 10 5s 2 5p 4 Polonium Po 84 [Xe] 4f 14 5d 10 6s 2 6p 4 General Characteristics: 1. Physical state: Oxygen is a gas at room temperature while other members of this group are solids. 2. Atomic radius: As the atomic number increases, the atomic radius increases from oxygen to tellurium. 3. Ionisation energy: Ionisation energy decreases gradually on descending the group from oxygen to polonium. 4. Electronegativity: Electronegativity decreases gradually on descending the group from oxygen to polonium. The gradual decrease in I.P. and E.N. is explained by the increase in size and screening or shielding ef fect. 5. Metallic character: As the atomic size increases, the metallic character increases from oxygen to polonium. O, S, Se and Te are non-metals. 6. Density : The density increase gradually from oxygen to tellurium. 7. Melting point : The melting point increases gradually from oxygen to tellurium. 8. Boiling point : The boiling point increases from oxygen to tellurium. The gradation in density, M.P and B.P is explained by the intermolecular attractive forces from oxygen to tellurium. 9. Atomicity : Oxygen molecule is diatomic sulphur, selenium and tellurium are octa atomic molecules. In oxygen molecule the two atoms are joined by a double bond. Sulphur molecule has a puckered ring structure. 10. Oxidation states: All the elements of this group except oxygen exhibit the principal oxidation states –2, +2, +4 and +6. The normal oxidation state of oxygen is –2. The very high electronegativity of oxygen suggests that it forms the compounds of –2 state. The electronegativities of other elements are, however low. If the chalcogen atom is more electronegative, in the molecule it shows – 2 oxidation state. If it is less electronegative it shows +2, +4, +6 oxidation states. Oxygen cannot form more than two bonds because of the non-availability of ‘d’ orbitals in its outershell. Much energy is required to excite the electrons to higher levels. The elements S, Se, Te and Po have vacant orbitals in the valence shell and can form four or six bonds in the excited state. 2, 4 or 6 bonds of sulphur can be explained as follows.

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Page 1: Oxygen Family - Theory Notes With Illustrative Examples [Unlocked by Www.freemypdf.com]

IIT-JEE Via Group Elements

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VIA GROUP ELEMENTS Introduction Oxygen, sulphur, selenium, Tellurium and polonium are the elements of VIA group. They have ns 2 np4 outer electron configuration. They belong to ‘p’ block in the long form of the periodic table. The first four elements are collectively known as Chalcogens (means mineral forming) since many metals occur as oxides or sulphides. Polonium is a radioactive element and short lived. It was first isolated by Madam Curie from an ore of Uranium called Pitchblende. Curie named this element as Polonium in honour of her native land, Poland. The elements of this group have six electrons in their outer shell. These electrons are described as valence electrons. In these six electrons two electrons are in ‘s’ orbital and four electrons are in ‘p’ orbitals. In these four electrons two are unpaired. These unpaired electrons generally take part in chemical bonding. All these elements require two electrons to attain a stable inert gas configuration.

Electron Configuration of VIA Group Elements Element Symbol Atomic Number Electron Configuration

Oxygen O 8 [He] 2s2 sp4 Sulphur S 16 [Ne] 3s2 3p4

Selenium Se 34 [Ar] 3d10 4s2 4p4 Tellurium Te 52 [Kr] 4d10 5s2 5p4 Polonium Po 84 [Xe] 4f14 5d10 6s2 6p4

General Characteristics: 1. Physical state: Oxygen is a gas at room temperature while other members of this group are solids. 2. Atomic radius: As the atomic number increases, the atomic radius increases from oxygen to tellurium. 3. Ionisation energy: Ionisation energy decreases gradually on descending the group from oxygen to polonium. 4. Electronegativity: Electronegativity decreases gradually on descending the group from oxygen to polonium.

The gradual decrease in I.P. and E.N. is explained by the increase in size and screening or shielding ef fect. 5. Metallic character: As the atomic size increases, the metallic character increases from oxygen to polonium. O,

S, Se and Te are non-metals. 6. Density : The density increase gradually from oxygen to tellurium. 7. Melting point : The melting point increases gradually from oxygen to tellurium. 8. Boiling point : The boiling point increases from oxygen to tellurium. The gradation in density, M.P and B.P is

explained by the intermolecular attractive forces from oxygen to tellurium. 9. Atomicity : Oxygen molecule is diatomic sulphur, selenium and tellurium are octa atomic molecules. In oxygen

molecule the two atoms are joined by a double bond. Sulphur molecule has a puckered ring structure. 10. Oxidation states: All the elements of this group except oxygen exhibit the principal oxidation states –2, +2, +4

and +6. The normal oxidation state of oxygen is –2. The very high electronegativity of oxygen suggests that it forms the compounds of –2 state.

The electronegativities of other elements are, however low. If the chalcogen atom is more electronegative, in the molecule it shows – 2 oxidation state. If it is less electronegative it shows +2, +4, +6 oxidation states. Oxygen cannot form more than two bonds because of the non-availability of ‘d’ orbitals in its outershell. Much energy is required to excite the electrons to higher levels. The elements S, Se, Te and Po have vacant orbitals in the valence shell and can form four or six bonds in the excited state. 2, 4 or 6 bonds of sulphur can be explained as follows.

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3s 3p 3d

Ground State

Two unpaired electrons can form two bonds. Due to sp3

3s 3p 3d

First Excited State

hybridisation, the central atom has a tetrahedral geometry with two positions occupied by lone pairs

Four unpaired electrons can form four bonds. Due to sp3

3s 3p 3d

Second Excited State

d hybridization, the atom acquires trigonal bipyramidal geometry with one position occupied by a lone pair.

Six unpaired electrons can to form six bonds. Due to sp3d2 hybridisation, the atom acquires octahedral geometry.

The higher oxidation state becomes less stable on descending the group. The +4 state shows both oxidizing and reducing properties while +6 state shows only oxidizing properties. These compounds are volatile because they are covalent.

Physi cal Properties of the VIA Group Elements Property Oxygen Sulphur Selenium Tellurium Polonium

Atomic weight 15.9 32.06 78.96 127.60 210.0 Atomic volume 14.00 15.50 16.50 20.50 22.7 Atomic radius (A) 0.73 1.09 1.16 1.35 -- Ionisation energy (K.Cal/mole)

314.00 253.90 231.00 199.10

184.90

Electronegativity 3.50 2.50 2.40 2.10 2.0 Density (g/c.c) 1.14 2.07 4.79 6.24 9.2 Melting point (K) 54.20 392.00 490.00 722.5 527.0 Boiling point (K) 90.0 717.60 958.00 1262.80 -- Oxidation states –2 –2, +2, +4,

+6 –2, +2, +4, +6

–2, +2, +4, +6

+2, +4,

11. Allotropy: All these elements exhibit allotropy. Oxygen exists in two non-metallic forms, oxygen (O 2) and ozone (O3

o

A

).

In ozone O – O bond length is intermediate between the single bond distance (1.48 ) and the double bond distance

(1.10 o

A ). It is 1.27 o

A . The actual structure of ozone is a resonance hybrid of two canonical structures I and II. Sulphur exists in many allotropic forms. All are non-metallic, e.g: Rhombic, monoclinic and plastic sulphur. Rhombic sulphur is stable at room temperature and it has a crystalline structure. Monococlinic sulphur is stable above 368.5K and it forms needle like crystals. The temperature 368.5K at which these two allotropes exist in equilibrium state is called transition temperature. These two forms are known as enantiotrophs. When sulphur is boiled above 473 K, S 8 rings are opened and long chains are formed. When this boiling sulphur is poured into water, a rubber like mass called plastic sulphur is obtained.

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ALLOTROPES OF SULPHUR

Allotrope State Specific Gravity M.P Structure Rhombic sulphur (α-sulphur)

Crystalline 2.06 385.8 Puckered ring

Monoclinic sulphur (β-sulphur)

Crystalline 1.98 392.2 Puckered ring

Plastic sulphur (γ-sulphur)

Amorphous - - Open chain S4

molecules Colloidal sulphur Colloid - - -

Selenium has two forms. They are red-non metallic and grey-metallic. Tellurium exists in two forms. They are non-metallic and grey-metallic. Polonium exists in two forms. They are α-polonium and β-polonium. Both are metallic.

Chemical Properties: 1. Hydrides : Chalcogens form the hydrides of the type H2X (where X is chalcogen). The affinity of these elements with hydrogen decreases from oxygen to polonium. So the stability of these hydrides decreases from H2O to H2Po. Actually hydrogen polonide is formed in traces. These hydrides are formed by the action of an acid on metallic sulphides, selenides, tellurides. FeS + H2SO4 → FeSO4 + H2S

Al2Se3 → + 6HCl 2AlCl3 + 3H2Se

Al2Te3 + 3H2SO4 → Al2Cl3 + 3H2Te

Hydrogen polonide is formed by the action of dilute acid on magnesium-polonium alloy. i. All these hydrides except water are gases at room temperature. Water is a liquid. ii. All these hydrides except water are volatile, poisonous and foul smelling gases, iii. water has an abnormally high melting and boiling points compared to the other hydrides.

This is due to inter molecular hydrogen bonding in water. iv. The volatility of these compounds decreases from H2S to H2Te. v. The reducing property of these hydrides increases from H 2O to H2Po. vi. These hydrides dissolve in water to give weak acids. H2X + H2 →O H3O+ + HX–

H2S + H2 →O H3O+ + HS–

In solution, the acidic property increases from H2S to H2Te. vii. All the molecules of these hydrides have bent structure. The bond angle decreases from H 2O to H2

Te.

O

H H 104°31’

H

H 92°

S

In the formation of water molecule, oxygen atom undergoes sp3 hybridisation. Because of the repulsion or non-bonding electron pairs, the tetrahedral bond angle is decreased. In other hydrides, the bond angle is close to 90 °. It shows that in other chalcogen atoms only pure valence p orbitals involve in the bond formation.

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Oxygen exhibits catenation to a small extent while sulphur exhibits to a greater extent. So they form polyoxides and polysulphides which are less stable than normal salts. e.g.: H2O2 : H – O – O – H; H2S2 : H – S – S – H; H2S3 : H – S – S – S – H When ice cold acid is added to a peroxide or persulphide, hydrogen peroxide or hydrogen persuilphide is formed. BaO2 → + 2HCl H2O2 + BaCl2

Na2S2 →+ 2HCl H2S2

Name of Hydride

+ 2NaCl

HYDRIDES VIA GROUP ELEMENTS

Formula Bond Angle Boiling Point Hydrogen oxide H2 HOH = 105 O 373 K Hydrogen sulphide H2 HSH = 92 S 213 K Hydrogen selenide H2 HSeH = 91 Se 231 K Hydrogen telluride H2 HTeH = 90 Te 270 K Hydrogen polonide H2 HPoH = 90 Po Unstable

2. Halides : The general formulae of VIA group halides are M2X2, MX2, MX4 and MX6 (M = S, Se and Te; X = halogen) in which the oxidation states of the chalcogen are +1, +2, +4 and +6 respectively. i. Oxygen combines with chlorine to form a number of oxides. These are chlorine monoxide (Cl 2O), chlorine

dioxide(ClO2), chlorine hexoxide (Cl2O6) and chlorine heptoxide (Cl2O7) in which the oxidation states of chlorine are +1, +4, +6 and +7 respectively.

ii. Other elements of this family combine with chlorine to form tetra chlorides. E.g: SCl 4, SeCl4, TeCl4 and PoCl4, SCl4 is very unstable liquid. The other tetra chlorides are stable solids. All these tetra chlorides undergo hydrolysis.

MCl4 + H2 O M(OH)4

+ 4 HCl

H2MO3 + H2O iii. Structure of these tetra halides is distorted trigonal bipyramid (trigonal bipyramid with one orbital occupied by

a lone pair of electrons). iv. The higher oxidation state of +6 is exhibited by the elements S, Se, and Te with fluorine. They form SF 6, SeF6

and TeF6. All these fluorides are formed by the direct combination of these elements with fluorine. v. These halides undergo hydrolysis and the degree of hydrolysis increases with the increasing atomic number.

Thus TeF6 hydrides TeF6 + 6H2 →O H6TeO6 + 6HF

vi. Structure of these fluorides is octahedron due to sp3d2

Element

hybridisation of the central atom. Halides of VIA group elements are listed in the table.

HALIDES OF VIA GROUP ELEMENTS

M2X MX2 MX2 MX4 XM6 Other Compounds 2

Oxygen O2F F2 2O

Br2O Cl2

-

O

- ClO2 BrO

O

2 3F2, O4F2, ClO6

Cl2O7, BrO3, I2O4, I4O9, I2O5

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Sulphur S2F2 S2Cl

SF

2 2

SClSF

2 4

SClSF

2

- 6 -

Selenium Se2Cl2 Se2Br

SeCl

2 SeF2 4 SeCl4 SeBr

SeF

4

- 6 -

Tellurium - TeCl2 TeBr

TeF

2 4

TeCl4 TeBr

TeF

4

- 6 -

Polonium - PoCl2 PoBr

PoCl

2 4

PoBr4 Pol

-

4

- -

3. Oxides i. All these elements combine with oxygen to form dioxides. MO 2 and trioxides MO3(M = S, Se, Te & Po) eg:

sulphur forms sulphur dioxide, SO2 and sulphur trioxide SO3. The dioxide is prepared by heating the element in air.

S + O2 → SO2

ii. As the metallic character increases the acidic nature of the oxides gradually decreases. iii. Reducing power gradually decreases and oxidizing power increases. iv. To prepare trioxides special methods are used. S, Se and Te form trioxides. SO 3, SeO3 and TeO3. Sulphur

trioxide is prepared by heating a mixture of SO2 and O2 over a catalyst. 2SO2 + O2

2 5V O→ 2SO3

v. The acidic nature of these oxides is in the order of SO 3 > SeO3 > TeO3 Tellurium and polonium give monoxides (TeO, PoO). 4. Oxyacids i. The dioxides of VIA group elements dissolve in water to form ‘ous’ acids of the type H 2MO3. The solubility of

dioxides decreases from SO2 to TeO2. MO2 + H2 →O H2MO3

SO2 + H2 →O H2SO3(sulphurous acid)

SeO2 + H2 →O H2SeO3(selenious acid)

TeO2 + H2 →O H2TeO3(tellurous acid)

PoO2 is insoluble in water. It is soluble onlyin acids. ii. The strength of the oxyacids is in the order of H2SO3 > H2SeO3 > Hs2TeO3 iii. In these acids, the oxidation state of VIA group element is +4. iv. The trioxides dissolve in water to form ‘ic’ acids of the type H 2MO4. MO3 + H2 →O H2MO4

SO3 + H2 →O H2SO4 (sulphuric acid)

SeO3 + H2 →O H2SeO4 (selenic acid)

TeO3 + H2 →O H2TeO4 (telluric acid)

v. Just as the strength of the ‘ous’ acids the strength of the ‘ic’ acids is in the order of H2SO4 > H2SeO4 > H2TeO4

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vi. In tyhese acids, the oxidation state of VIA Group element is +6. Among the oxyacids of VIA group elements, oxyacids of sulphur are important (shows in the table)

Oxy Acids of Sulphur

Name Formula Structure Oxidation state

Sulphoxylic acid H2SO HO – S – OH 2 + 1

Sulphurous acid H2SOO = S – OH

| OH

3 + 4

Di or pyrousulphorus acid H2S2O

O ||

HO – S – S – S – OH || || O O

5 + 4

Sulphuric acid H2SO

O ||

HO – S – OH || O

4 + 6

Thiosulphuric acid H2S2O

S ||

HO – S – OH || O

3 + 2

Pyrosulphuric acid H2S2O

O O || ||

HO – S – O – S – O H || || O O

7 + 6

Dithionic acid H2S2O

O O || ||

HO – S – S – O H || || O O

6 + 5

Peroxy monosulphuria acid (or) permonosulphuric acid (or) Caro’s acid

H2SO

O ||

HO – O – S – O H || O

5 + 6

Peroxo disulphuric acid (or) perdisulphuric acid (or) Marshall’s acid

H2S2O

O O || || HO – S – O – O – S – O H

|| || O O

8 + 6

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Ozone, O 3

2 33O 2O 68 K. Cal.−

. Introduction: Ozone was first noticed by Van Marum in 1785. Schenbien named it as Ozone because of its peculiar smell. The triatomic nature was observed Soret in 1866. He considered it as an allotrope of oxygen. Ozone is largely present in the upper layer of the atmosphere. It prevents the powerful ultra-violet radiation from the Sun. PREPARATION OF OZONE Ozone is generally prepared by subjecting oxygen to silent electric discharge in an apparatus called Ozonizer There are two important ozonizers. They are: 1. Siemen’s Ozonizer and 2. Brodie’s Ozonizer Siemen’s Ozonizer: Siemen’s ozonizer consists of two coaxial glass tubes sealed at one end. The inner surface of the inner tube and outer surface of the outer tube are coated with tin foil. These tin foils are connected to the terminals of a powerful induction coil. A current of pure and dry oxygen gas is passed through the annular space and is subjected to silent electric discharge. Ozonized oxygen which is a mixture of oxygen with 3 to 8% ozone comes out from the other end.

2. Brodie’s ozonizer: Brodie’s ozonizer consists of two concentric glass tubes. The inner tube contains dilute sulphuric acid and the outer tube is suspended in dilute sulphuric acid taken in a beaker. Two copper wires form the electrodes. One electrode is dipped in dilute sulphuric acid of the inner tube and the other in dilute sulphuric acid in the beaker. These wires are connected to a powerful induction coil. A current of pure and dry oxygen gas passed through the annular space and subjected silent electric discharge. Ozonized oxygen comes out through the outlet.

2 33O 2O 68 K. Cal−

MANUFACTURE OF OZONE Ozone is manufactured by Siemen – Halske’s ozonizer. It consists of about six to eight porcelain or glass cylinders. Each is surrounded by a cylinder of aluminium fitted in a metal box. These aluminium cylinders are kept in vertical position on an insulating glass plate fitted into an iron tank. This tank is divided into three compartments. Through these compartments cold water is circulated to keep the apparatus cool. The aluminium rods are raised to a high potential of 8000 – 10000 volts. Air is passed through the annular space around the aluminium cylinders. Then the air is subjected to the action of silent electric discharge and comes out in the form of ozonized air.

2 33O 2O 68 K. Cal−

PHYSICAL PROPERTIES 1. Ozone is a pale blue gas at room temperature. 2. It has a characteristic fishy odour. 3. In liquid state ozone is blue in colour.

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4. It is slightly more soluble in water than oxygen. It is soluble in turpentine oil. 5. It acts as a respiratory irritant if present in air in larger proportions. CHEMICAL PROPERTIES 1. Ozone slowly decomposes into oxygen at room temperature but rapidly at 473 K. The rate of decomposition is

increased in presence of catalysts like silver, platinum, vanadium etc.

3 22O 3O 68 K. Cal.→ +

2 vol 3 vol 2. Oxidising reactions: Ozone acts as a powerful oxidizing agent. Usually only one atoms of oxygen from each

molecule is available for bringing about oxidation. Thus the principal equation is :

3 2O O O→ +

i. Ozone oxidizes lead sulphide (black) to lead sulphate (white)

3 2

4

3 4 2

O O O 4

PbS 4O PbSO

PbS 4O PbSO 4O

→ + × + →+ → +

ii. It oxidizes hydrohalic acids to halogens.

3 2

4

3 2 2 2

O O O

2HCl O PbSO

2HCl O H O Cl O

→ ++ →+ → + +

iii. It oxidizes potassium maganate solution (green) to potassium permanganate solution (pink).

3 2

2 4 2 4

2 4 2 3 4 2

O O O

2K MnO H O O 2KMnO 2KOH

2K MnO H O O 2KMnO 2KOH O

→ ++ + → ++ + → + +

iv. it oxidizes potassium ferrocyanide solution to potassium ferricyanide solution.

3 2

4 6 2 3 6

4 6 2 3 3 6 2

O O O

2K Fe(CN) H O O 2K Fe(CN) 2KOH

2K Fe(CN) H O O 2K Fe(CN) 2KOH O

→ ++ + → ++ + → + +

v. Ozone liberates iodine from moist potassium iodide.

3 2

2 2

2 3 2 2

O O O

2KI H O O 2KOH I

2KI H O O 2KOH I O

→ ++ + → ++ + → + +

vi. Ozone directly oxidizes moist iodine, sulphur, phosphorus, arsenic, antimony etc. to their corresponding oxyacids with highest oxidation state.

I2 + H2O + 5O2 → 2HIO3 + 5O2

S + H2O + 3O3 → H2SO4 + 3O2

P4 + 6H2O + 10O2 → 4H3PO4 + 10O2

2As + 3H2O + 5O3 → 2H3AsO4 + 5O2

2Sb + 3H2O + 5O3 → 2H3 SbO4 + 5O2

vii. Tailing of mercury: Mercury reacts with ozone to form mercurous oxide 2Hg + O2 → Hg2O + O2

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The mercurous oxide produced dissolves in mercury which loses its mensiscus and sticks to the walls of the glass. This phenomenon is known as tailing of mercury. When this mercury is washed with water, it restores its original properties. viii. Exceptional oxidation reactions: In few cases, all the three atoms of ozone molecules are used up for

oxidation. 3SO2 + O3 → 3SO3

3SnCl2 + 6HCl + O3 → 3SnCl4 + 3H2O

3. Reducing properties: In some reactions ozone acts as a reducing agent. Ozone takes up an oxygen atom from the substance for bringing about reduction. Thus, the principal equation is: O3 → + O 2O2

i. Ozone reduces barium peroxide to barium oxide. BaO2 + O3 → BaO + 2O2

ii. It reduces hydrogen peroxide to water.

H2O2 + O3 → H2O + 2O2

iii. It reduces silver oxide to metallic silver. Ag2O + O3 → 2Ag + 2O

CH

2

4. Bleaching action Ozone combines with unsaturated hydrocarbons to form very unstable compounds called ozonides. These products are hydrolysed to carbonyl compounds and hydrogen peroxide. This process is known as ozonolysis. It is used to locate the position of carbon-carbon multiple in the original unsaturated compound. E.g:

2 = CH2 + O3 →

Ethylene

O

O O

H2C CH2

ethyleneozonide

O

O O

H2C CH2 + H2O

→ 2HCHO + H2O

ii. HC = CH + O

2

formaldehyde

3 →

O

O O

HC CH +H2O CHO

CHO glyoxal

+ H2O2

Uses of Ozone Ozone is used 1. as a germicide and disinfectant for purification of air and sterilization of water. 2. as a mild bleaching agent for ivory, flour, starch, oils, wax etc. 3. to detect double and triple bonds in organic compounds. 4. to purify air in underground railways, zoos, mines, cinema halls etc.

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5. a mixture of ozone and cyanogens (C2N2

) is used as rocket fuel. Tests of Ozone 1. It has a strong fishy like odour. 2. It turns starch iodide paper blue. 3. It turns benzidine paper brown and tetramethyl base paper violet. 4. Tailing of mercury: Mercury loses its meniscus in contact with ozone and sticks to the surface of glass. 5. A clean silver foil is blackened by ozone. Formula of Ozone i. Ozone is prepared by subjecting oxygen to silent electric discharge. It gives oxygen on heating. This shows

that it is made up of oxygen atoms only. ii. Soret’s experiment: Soret took two graduated flasks of exactly same capacity. These flasks were filled with

exactly the same sample of ozone. These flasks were inverted in a trough of water. In one of the flasks (A) he introduced turpentine oil and heated the other (B). A decrease in volume was noticed in the flask A due to the absorption of ozone. An increase in volume was noticed in the flask B due to the decomposition of ozone to oxygen. The decrease in volume was found to be double the increase i.e. If the increase in volume is one, the decrease is two volumes. The decrease in volume (2 volumes) is the volume of ozone present in ozonised air. This increases by one volume on heating i.e. when 2 volumes of ozone decomposes, 3 volumes of oxygen is obtained. Ozone Oxygen

2 vol 3 vol Applying Avogadro’s law 2n molecules of ozone yield 3n molecules of oxygen. Two molecules of ozone yield 3 molecules of oxygen. One molecule of ozone yields 3/2 molecules or 3 atoms of oxygen. Hence formula of ozone should be O 3

o

A

. Structure of Ozone:

Ozone is diamagnetic and has an angular structure. Both the oxygen to oxygen bonds have the same length (1.27 )

which is intermediate between the double bond distance (1.10 o

A ) and a single bond distance (1.48 o

A ). The molecule may be represented as a resonance hybrid of the following structures.

O

.. O O :

..

: : ..

O

.. O: O

..

: : ..

+ +

– –

Spectroscopic studies suggest a non-linear structure as shown below.

O

O O 117°

Manufacture of Sulphuric Acid, H 2SO4: It is prepared on a large scale by two processes. 1. Lead chamber process and 2. Contact process

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1. Lead chamber process: It is an older method for the manufacture of sulphuric acid. In this process the chambers made up of lead are used. So this is called Lead chamber process. The metal lead is used because it is cheaper and has no action with sulphuric acid. Principle: In this process SO2 is oxidized catalytically by means of oxides of nitrogen in the presence of water to form

nitroso sulphuric acid. N2O3 + O2 + H2O + 2SO2 → 2 NOHSO4

When nitroso sulphuric acid is hydrolysed sulphuric acid is formed. 2NOHSO4 + H2 →O 2H2SO4 + NO + NO2

The oxides of nitrogen simply act as carriers of oxygen from air to sulphur dioxide. This method involves homogenous catalysis. Conditions: The above reactions take place at 673 to 733 K and 1.5 – 2 atmospheres pressure. 2. Contact Process: Principle: Purified SO2 is oxidized to SO3 catalytically. SO3 vapour is absorbed in concentrated H2SO4 to form pyrosulphuric acid or Oleum, H2S2O7. Water is added to oleum to form sulphuric acid of the desired concentration. 2 SO2 + O2 → 2 SO3

SO3 + H2SO4 → H2S2O7

H2S2O7 + H2 →O 2 H2SO4

Conditions: Oxidation of sulphur dioxide is reversible, exothermic and it proceeds with decrease in volume. According Le Chatelier’s principle, the optimum conditions for the better yield are: i. Low temperature: The oxidation of sulphur dioxide to sulphur trioxide is exothermic. So this is favoured by low

temperature. The reaction is too slow at low temperature. Hence the oxidation is carried out at about 723 K. ii. High pressure: High pressure favours the oxidation. However, at high pressure, the material of the plant

corrodes. So the reaction is carried out at a pressure slightly higher than one atmosphere by using excess of oxygen.

iii. The presence of catalyst in contact process lowers the activation energy of the reacting gases. It also increases the number of fruitful collisions. These factors favour the formation of sulphurtrioxide. Powdered state of the catalyst is more effective. Vanadium pentoxide is used as a catalyst and a compound of potassium is uses as a promoter.

The plant used for the manufacture of sulphuric acid by contact process is mainly consists of four units. 1. Sulphur or pyrites burners: In pyrites burners sulphur dioxide is obtained either by burning sulphur in air or by

roasting iron pyrites. S + O2 → SO2

4FeS2 + 11O2 → 2Fe2O3 + 8SO2

The gases contain impurities like arsenic trioxide, sulphur or pyrite dust and sulphuric acid fog. Thy must be removed to prevent the poisoning of the catalyst.

2. Purification unit: It consists fo the following parts i. Dusting tower or precipitator: In this tower steam is blown in, to make the dust particles settle down.

Sometimes dust particles are precipitated mechanically or electrically in precipitator. ii. Cooling pipe: The gases are cooled by passing them through pipe, cooled by air. The gases moving upwards

are washed by a stream of water flowing down. iii. Drying tower: The moist gases are dried in drying tower by a spray of concentrated sulphuric acid. iv. Arsenic purifier: The dried gases are passed through arsenic purifier to be free from arsenic impurit ies. It

contains gelatinous ferric hydroxide.

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v. Testing box: The gases are passed into a testing box to find whether thy are free from impurities or not. In the box, a strong beam of light is passed against the gases. If any particles are present, they get illuminated due to Tyndall effect.

3. Contact chamber: Pure gases are passed through a preheater and is heated to about 723K. These hot gases

are passed into the contact chamber. In contact chamber sulphur dioxide and oxygen combine in presence of a catalyst vanadium pentoxide to form sulphur trioxide.

2SO2 + O2 2SO3 + 45.2K.Cal

The reaction is exothermic and the heat produced raises the temperature of the catalyst to about 723 K. Once the reaction starts, the gases need not be heated in heat exchanger.

4. Absorption tower: The gases containing mainly sulphur trioxide and nitrogen enter the absorption tower. Here

sulphur trioxide is absorbed by 98% sulphuric acid to give oleum. SO3 + H2SO4 → H2S2O7

Sulphuric acid of desired strength can be obtained by diluting the oleum with water. H2S2O7 + H2 →O 2H2SO4

Properties of sulphuric acid Physical properties 1. Pure sulphuric acid is a colourless syrupy liquid. 2. Its density is 1.84 g/ml. 3. Sulphuric acid forms hydrates with explosive violence. So the affinity for water is very high. Chemcial Properties 1. Acidic nature: Sulphuric acid is a very strong dibasic acid. It reacts with alkalies forming two types of salts,

bisulphates and sulphates. NaOH + H2SO4 → NaHSO4 + H2O

2 NaOH + H2SO4 → Na2SO4 + 2H2O

2. Action on non-metals: Sulphuric acid oxidizes carbon to carbondioxide, sulphur to sulphurdioxide and

phosphorus to phosphoric acid. C + 2H2SO4 → 2H2O + 2SO2 + CO2

S + 2H2SO4 → 2H2O + 3SO2

2P + 5H2SO4 → 2H3PO4 + 2H2O + 5SO2

3. Action on metals: Metals like copper, silver, mercury, antimony which lie below hydrogen in electrochemical series do not evolve hydrogen with dilute sulphuric acid. These metals react in presence of air or oxygen to form sulphates and water

2Cu + 2H2SO4 + O2 → 2CuSO4 + 2H2O

Concentrated sulphuric acid reacts with these metals on boiling and liberate sulphur dioxide. Cu + 2 H2SO4 → CuSO4 + 2H2O + SO2

2Ag + 2 H2SO4 → Ag2SO4 + 2H2O + SO2

Metals like Zn, Mg, Na, K, Al, Pb, Sn etc., which lie above hydrogen in electrochemical series react with dilute

H2SO4 to liberate H2. Zn + H2SO4 → ZnSO4 + H2

Mg + H2SO4 → MgSO4 + H2

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These metals react with concentrated sulphuric acid and liberate sulphurdioxide. Zn + H2SO4 → ZnSO4 + 2H2O + SO2

2Al + 6H2SO4 → Al2(SO4)3 + 6H2O + 3SO2

Metals like lead, gold, platinum do not react with sulphuric acid. 4. Action with halides: Concentrated sulphuric acid reacts with fluorides and chlorides to form hydrogen fluoride

and hydrogen chloride respectively. CaF2 + H2SO4 → CaSO4 + 2HF

NaCl + H2SO4 → NaHSO4 + HCl

But hydrogen bromide and hydrogen iodide which are formed from bromides and iodides are oxidized to bromine and iodine respectively.

2KBr + H2SO4 → K2SO4 + 2HBr

2HBr + H2SO4 → 2H2O + SO2 + Br2

2KI + H2SO4 → K2SO4 + 2HI

2HI + H2SO4 → 2H2O + SO2 + I2

These reactions are used for the detection of halides. 5. Affinity for water: It has great affinity for water. Therefore it is used for drying all gases except ammonia and

hydrogen sulphide. It is also used as a strong dehydrating agent. a. Paper, starch, wood etc, are all charred by concentrated sulphuric acid due the removal of water. It is

corrosive towards skin and causes painful blisters. b. Sugar is charred to give charcoal, Charring of sugar is due to dehydration. C2H22O11

2 4conc.H SO→ 12C + 11H2

2 4conc.H SO2 2

COOH| H O CO COCOOH

→ + +

O

c. It removes water from oxalic acid and formic acid.

2 4conc.H SO2HCOOH H O CO→ +

d. Conc.H2SO4

o170 C2 2 2 4 2C H OH C H H O→ +

removes water from alcohol to give ethylene.

Uses of Sulphuric Acid: It is used 1. as a reagent in the laboratory. 2. as a dehydrating and drying agetnt. 3. in fertilizer industry for the manufacture of ammonium sulphate and superphosphate of lime. 4. in refining petroleum and in coaltar industry. 5. in paper, textile, leather and rubber industries. 6. in the manufacture of dyes, drugs and disinfectants. 7. in the manufacture of paints and pigments.

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8. in the manufacture of important chemicals like hydrochloric acid, nitric acid, phosphoric acid, alums, sulphates, ether etc.

9. in the manufacture of explosives like T.N.T dynamite, nitroglycerine etc. 10. in storage batteries. 11. in metallurgy. Sodium Thiosulphate (Na 2S2O3.5H2O) Sodium thiosulphate is also known as Hypo. Preparation: Sodium thiosulphate is prepared by boiling sodium sulphite solution with sulphur. Na2SO3 → + S Na2S2O3

The excess of sulphur is filtered off and the filtrate is evapourated to form sodium thiosulphate crystals, Na 2 S2 O3. 5H2O

Properties: 1. It is a colourless crystalline solid. 2. It is highly soluble in water. 3. Action of heat: On heating thiosulphate loses water of crystallization at 488 K. Na2S2O3.5H2 →O Na2S2O3 + 5H2O

It decomposes when heated above 493 K to form sodium sulphate and sodium penta sulphide. 4Na2S2O3 → 3Na2SO4 + Na2S5

4. Action of dilute acids: Dilute acids decompose thiosulphate with the liberation of sulphur dioxide an d precipitation of sulphur.

Na2S2O3 →+ 2HCl 2NaCl + H2O + S + SO2

5. Action of silver nitrate: Silver nitrate reacts with hypo to form a white precipitate which quickly changes to different colours and finally to black.

Na2S2O3 + 2AgNO3 → Ag2S2O3 + 2NaNO3

White Ag2S2O3 + H2 →O Ag2S + H2SO4

black 6. Action of ferric chloride: Ferric chloride reacts with hypo to form a violet colour ferric thiosulphate.

3Na2S2O3 + FeCl3 → Na3[Fe(S2O3)3] + 3NaCl

7. Action with iodine solution: Sodium thiosulphate decolourises iodine solution to form tetrathionate. 2Na2S2O3 + I2 → 2NaI + Na2S4O6

This reaction is used in estimation of iodine in volumetric analysis. 8. Action of chlorine: Chlorine oxidizes sodium thiosulphate solution to sodium sulphate. Na2S2O3 + H2O + Cl2 → Na2SO4 + S + 2HCl

Hence hypo is used as antichlor in textile industry. 9. Action of silver halides: Silver halides are easily soluble in hypo solution. So hypo is used in photography as

developer. AgX + 2Na2S2O3 → Na3[Ag(S2O3)2] + NaX

Uses of Sodium Thiosulphate: It is used 1. in photography for fixing. 2. in textile industry as an antichlor. 3. in the extraction of silver and gold. 4. in the iodometric titrations. 5. in medicine.

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