Chemistry Book Two

ABDULLAH EDUCATIONAL SOCIETY 110/2, Gulbahar No. 2

CHEMISTRY

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Contents HYDROGEN AND WATER 3

Q - 1) Define Hydrogen and write down uses, physical and chemical properties of Hydrogen: ............................... 3

Q - 2) What is nascent hydrogen? Describe its reactivity: .............................................................................................................. 5

Q - 3) Define isotope and discus various types of hydrogen isotopes: .................................................................................... 5

Q - 4) How is hydrogen prepared commercially: ............................................................................................................................... 6

Q - 5) Give reaction H2 with: ....................................................................................................................................................................... 7

Q - 6) Define Water and describe some main physical and chemical properties of Water: ........................................ 8

Q - 7) What do you understand by the anomalous behaviour of water? And what is the significant of this

unusual behaviour of water? ......................................................................................................................................................... 9

Q - 8) What is water of crystallization? What happens when hydrates are heated? Define heat of hydration also

write formulas of some of hydrates. ........................................................................................................................................ 10

Q - 9) What is “Potable water”? Write four main characteristics of it. ..................................................................................... 10

Q - 10) What do you mean by soft water, hard water and heavy water? ................................................................................. 10

Q - 11) What do you mean by hard water and describes the types of hardness? Also write causes of hardness... 11

Q - 12) How is the hardness of water removed? Describe various methods to remove hardness of water? ............ 12

Q - 13) Describe disadvantages of hard water and uses of water? ............................................................................................. 13

Q - 14) Describe the classification of water Pollutants: or name water borne diseases that are caused by

microorganism presents in water, name various types of water pollutants and their different categories.

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Q - 15) Describe hygroscopic substances: ............................................................................................................................................ 14

Q - 16) Name only some common treatment to make municipal water fit for drinking purposes. .............................. 15

Q - 17) What happens when? ..................................................................................................................................................................... 15

CARBON, SILICON AND THEIR COMPOUNDS 16

Q - 1) Describe Carbon and its uses. Also write its physical and chemical properties: .................................................... 16

Q - 2) Define allotropy and allotropes. Discuss chief allotropic forms of carbon, also write their uses & structure.

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Q - 3) Discus amorphous forms of Carbon. Or Short note .......................................................................................................... 20

Q - 4) Define catenation: ........................................................................................................................................................................... 20

Q - 5) What is silicon? How does silicon occur in nature? ........................................................................................................... 21

Q - 6) What are silicates? And Describe some common silicates and their chemical formulas and uses. ............... 21

Q - 7) Give some physical properties of silicon. ............................................................................................................................... 22

Q - 8) Describe chemical preparation of Silicon. .............................................................................................................................. 22

Q - 9) Give some uses of silicon: ............................................................................................................................................................ 23

Q - 10) Describe preparation of Silica and its properties and uses:............................................................................................ 23

Q - 11) Describe preparation of Sodium Silicate or water glass and its properties and uses: ......................................... 24

Q - 12) What is Silica Gel? Also write some uses of it. ..................................................................................................................... 24

NITROGEN AND OXYGEN 25

Q - 1) Describe Nitrogen, its occurrence and uses. Also write physical and chemical properties: .............................. 25

Q - 2) Describe Oxygen and its occurrence & uses. Also write its physical and chemical properties: ....................... 26

Q - 3) Describe chemical preparation of Nitrogen. How can you get nitrogen from the atmospheric air? Give

tow uses of nitrogen. ...................................................................................................................................................................... 27

Q - 4) Describe laboratory preparation of Oxygen. How oxygen is industrially produced from liquid air? ............. 28

Q - 5) What are oxides how are they classified describe normal oxides in detail? ............................................................ 28

Q - 6) Give the preparation and properties of hydrogen peroxide and its uses: ................................................................ 30

Q - 7) Define oxidation and reduction and write chemical equation in support of each. ............................................... 32

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Q - 8) Define oxidizing agent, reducing agent and redox reactions. ....................................................................................... 33

Q - 9) What is ozone? How ozone is produced in the atmosphere? Or how oxygen is converted into ozone? Its

physical/chemical properties and what is important of ozone and write uses of ozone. .................................. 33

Q - 10) What is Aqua Regia? How does it dissolve gold? ............................................................................................................... 34

Q - 11) Give the preparation and properties of Ammonia (NH3) and its uses: ...................................................................... 35

Q - 12) Give the preparation, properties of Nitric Acid and its uses: ......................................................................................... 37

SULPHUR AND ITS COMPOUNDS 40

Q - 1) Describe Sulphur and its occurrence, uses, its physical and chemical properties: ................................................ 40

Q - 2) Describe and explain how Sulphur is extracted from underground deposits by Frasch process: ................... 41

Q - 3) Describe Sulphuric Acid and its occurrence, uses, its physical and chemical properties: ................................... 42

Q - 4) Describe and explain how Sulphuric Acid is prepared by Contact process, what catalyst used in the

process? ............................................................................................................................................................................................... 44

Q - 5) What is allotropy? Describe different allotropic forms of Sulphur? ............................................................................ 46

Q - 6) What is plastic Sulphur? Why is it elastic? ............................................................................................................................. 47

Q - 7) Explain why SO2 is purified before it is passed into the contact tower for its oxidation to SO3? .................... 47

Q - 8) Describe what happens when: .................................................................................................................................................... 47

HALOGENS 48

Q - 1) What are Halogens? Why are they placed in VII A group in the periodic table? Describe the

state of each member of the family and also their colours? .............................................................................. 48

Q - 2) What are the sources of halogens? Describe the importance of Cl2, Br2, and I2, F2 in our daily life.

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Q - 3) (a) How is chlorine prepared in the laboratory? (b) Describe the commercial preparation of

chlorine by the electrolysis of aqueous NaCl solution in Nelson’s Cell (c) what happens when

chlorine reacts with (i) Sn (ii) H2S (iii) CO (iv) P (v) FeCl2 (vi) H2O Give reaction and equation (d)

Discuss the uses of chlorine. .................................................................................................................................................. 50

Q - 4) Give the preparation of hydrogen chloride (HCl) in the laboratory by the action of conc. H2SO4

over common salt. (b) Describe the commercial preparation of Hydrogen chloride by the direct

combination of H2 and Cl2 gases (c) what happens when Hydrogen chlorine reacts with (i)

NaOH (ii) NaOHCO3 (iii) Pb(NO3)2 (iv) MnO2 (v) KMnO4 Give reaction and equation. (d) Discuss

the uses of HCl............................................................................................................................................................................... 55

Q - 5) What is bleaching powder? Give the preparation of bleaching powder by Hasenclever process

(b) Describe physical and chemical properties (c) what happens when bleaching powder reacts

with (i) water (ii) excess of water (iii) H2CO3 (weak acid) (iv) Ammonia (NH3) (v) HCl (d) Discuss

the uses of Bleaching powder. ............................................................................................................................................. 57

Q - 6) Describe Silver Nitrate (AgNO3) test for the presence of Cl- ion in a salt of chlorine. Give the

reaction. ............................................................................................................................................................................................. 59

Q - 7) Identify the following: ................................................................................................................................................................ 59

METALS AND THEIR EXTRATION 60

Q - 1) What do you mean by metal and non-metal? Also describe the physical and chemical properties of

metals and non-metals. ................................................................................................................................................................. 60

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Q - 1) Define Hydrogen and write down uses, physical and chemical properties of Hydrogen:

Hydrogen: A colourless, highly flammable gaseous element, the lightest of all gases and the most abundant element in the universe, used in the production of synthetic ammonia and methanol, in petroleum refining, in the hydrogenation of organic ma-terials, as a reducing atmosphere, in ox hydrogen torches, and in rocket fuels. Atomic number 1; atomic weight 1.00794; melting point -259.14°C; boiling point -252.8 °C; density at 0 °C 0.08987 gram per litter; valence 1. Physical Properties of Hydrogen: 1. It is a colourless, odourless and tasteless gas. 2. It is insoluble in water and highly inflammable gas and burns with blue flame. 3. Its bond dissociation energy is 104 k. cals/mol and ionization energy is 13.54 eV. 4. It liquefies at -252 °C and freezes at -259 °C. 5. Electro negativity of hydrogen is 2.1. Uses of Hydrogen: 1. It is used in manufacture of fertilizers. 2. It is used in manufacture of tungsten bold filaments and vegetable ghee. 3. It is used for purification of metals. 4. It is used in the preparation of oxy-hydrogen flame, which is used in welding due to production

of high temperature. 5. It is used in the preparation of chemicals like NH2, CH3OH, etc. Chemical Properties of Hydrogen:

1. Decomposition of molecular Hydrogen (H2) Molecular hydrogen contains stable covalent bonding and is relatively inert at ordinary conditions. Its bond dissociation energy is 104 K.cals/mol.

H — H H + H H = 104 K.cals/mol. i.e +435 KJ/,mol.

Q. Show hydrogen is a good reducing agent: 2. As a reducing Agent

Hydrogen shows greater affinity for oxygen and reduces many metal oxides into free metals. CuO(s) + H2(g) Cu(s) +3H2O(g)

WO3(s) + 3H2(g) W(s) + 3H2O(g)

CHEMISTRY HYDROGEN AND WATER

CHAPTER

IX ELEVENTH

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3. Hydrogenation Reaction: The addition of hydrogen into other molecular compounds called hydrogenation reaction. When molecular compounds and Hydrogen are heated in the presence of Pt or Pd or Ni and other crystals they give addition products.

ZnO/Cr2O3 1. CO(g) + 2H2(g) CH3 – OH(l)

400oC /high presusure Methyl Alcohol

Ni 2. CH2 = CH2(s) + H2(g) CH3 – CH3

Ethene 400oC Ethane

Ni 3. Edible oils (liquid)+ H2(s) Vegetable Ghee (Solid) Un saturated High Temp. Saturated

4. Reaction with metals: Alkali metals like Na, K etc. and alkaline earth metals like Ca, Ba react with hydrogen on heating to form ionic hydrides.

200 oC

1. 2Na(s) + H2(g) 2Na+H –(s) Sodium Hydride 200 oC

2. Ca(s) + H2(g) Ca+2 H2(s)–

Calcium hydride

5. Reaction with non-metals: Hydrogen reacts with many non-metals under different conditions to form addition products.

500 oC/200dtm 1. N2(g) + 3H2(g) 2NH3 (g)

Fe2O3/K2O

Sunlight 2. H2(g) +Cl2(g) 2HCl(g)

450 oC 3. H2(g) +S(s) H2S(g)

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Q - 2) What is nascent hydrogen? Describe its reactivity:

Nascent Hydrogen: Hydrogen at a time of its birth is chemically more reactive than molecular hydrogen because it is available in atomic form and known as Nascent Hydrogen or Newly born hydrogen. Reactivity: When a piece of Zn metal adds in the acidic Ferric chloride (FeCl3) solution, nascent hydrogen is generated which reduces FeCl3 into ferrous chloride (FeCl2) which is greenish in colour.

FeCl3(aq) + H2(g) No reaction Zn/HCl

FeCl3(aq) + [H] FeCl2(aq) + HCl(aq)

Nascent Hy-drogen

Ferrous Chloride

Acidic KMnO4 (pink) solution can be reduced by nascent hydrogen to colourless solution.

Nascent Hydrogen Zn/H2SO4

2KMnO44(aq) + 3H2SO4 + 10[H] K2SO4(aq) + 2MnSO4(aq) +8H2O Q - 3) Define isotope and discus various types of hydrogen isotopes:

Isotope: Isotope is the same atom with same number of protons but different num-ber of neutrons in its respective nuclei. Isotopes of Hydrogen:

There are three isotopes of hydrogen namely Protium, Deuterium and Tritium. 1. Protium or Ordinary Hydrogen Atom 𝐇𝟏

𝟏 :

This isotope contains one proton and one electron in the first orbit.

Its atomic number is 1 and mass number is also 1.

About 99.98% of free hydrogen contains Protium.

It is stable isotope of hydrogen.

2. Deuterium (D or 𝐇𝟏𝟐 ):

This isotope contains one proton, one neutron and one electron in the first orbit. Its atomic number is 1 and mass number is 2.

About 0.0156% of free hydrogen contains Deuterium.

It is also stable and heavy isotope of hydrogen i.e. misnomer.

3. Tritium (T or 𝐇𝟏𝟑 ):

This isotope contains one proton, two neutrons and one electron in the first orbit. Its atomic number is 1 and mass number is 3.

About 4 x 10–15 % of free hydrogen contains Tritium.

It is radioactive isotope with half-life of about 12.5 years.

It is used as tracer in the nuclear reactions.

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Q - 4) How is hydrogen prepared commercially:

Industrial Preparation of Hydrogen: Hydrogen is prepared by the following meth-ods. (i) By Passing steam over coke (ii) From Natural Gas (iii) By Thermal decompo-sition of methane (iv) By the electrolysis of water (i) By passing steam over coke: (Coke-steam process) When steam passes over red-hot coke at about 1000 oC, a mixture of carbon monoxide (CO) and hydrogen (H2) (Water Gas) produced.

1000 oC

C(s) + H2O(g) CO(g) + H2(g) ⏟

𝑊𝑎𝑡𝑒𝑟 𝐺𝑎𝑠

Water gas is a very good fuel and used in preparation of methanol (Methyl Alcohol). Hydrogen (H2) is separated by two methods from water gas: (i) By liquefaction (ii) By Oxidation

1. Separation of H2 by Liquefaction: When water is cooled up to – 200 oC, carbon monoxide liquefies and leaving behind H2 Gas. If the traces of CO gas left in the mixture, than the remaining mixture treats with caustic soda (NaOH) to form sodium formate and leaving behind pure H2 gas.

CO(g) + NaOH(aq) HCOONa(aq)

Carbon monoxide

Caustic Soda Solution

𝑆𝑜𝑑𝑖𝑢𝑚 𝐹𝑜𝑟𝑚𝑎𝑡𝑒

Q. Give Bosch method to separate hydrogen gas from water gas

2. Separation of H2 by Oxidation (Bosch Process): This is the most suitable method for the separation of H2 gas from water gas, in this process more steam passes through water gas at 500 oC in the presence of iron oxide (FeO) or chromium oxide (Cr2O3) catalyst. Carbon monoxide gas in water gas oxidizes to CO2 gas, which is soluble in water under pressure, lib-erating H2 gas.

FeO

CO(g) + H2(g) ⏟ + H2O(g) CO2(g) + 2H2(g)

Water gas Water 500 oC Solouble in water

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(ii) By Natural Gas: (Hydrocarbon-steam process) When steam passes over hydrocarbon such as methane (CH4) which is major constituent of natural gas in the presence of Nickel at about 900 oC, a mixture of carbon monoxide (CO) and hydrogen (H2) (Water Gas) produced.

900 oC

CH4(g) + H2O(g) CO(g) + 3H2(g) ⏟

Ni 𝑊𝑎𝑡𝑒𝑟 𝐺𝑎𝑠 (iii) By Thermal Decomposition of Methane: Hydrogen Gas (H2) is also commercially prepared by thermal decomposition of methane (CH4), when methane (CH4) is heated in the absence of Air at about 700 oC, methane (CH4) decomposes thermally to produce carbon black C(s) and H2 gas.

Above 700 oC

CH4(g) C(s) + 2H2(g) Absence of air Carbon Black

Uses of Carbon Black C(S): It is used in rubber industry as filler for manufacturing motor tyers. It is also used in the preparation of ink, paints, polishers, carbon papers and plas-

tics. (iv) By The Electrolysis of Water: Hydrogen Gas (H2) can also be produced by the electrolysis of water. When electric current is passed through water in the presence of a few drops of acid or base, H2 gas is liberated at cathode and O2 gas a by – product collects at anode. The presence of acid or base helps in the ionization of water.

Electricity

H2O(l) 2H2(g) + O2(g) Acid or Base Hydrogen gas

Q - 5) Give reaction H2 with: (i) Ethene (ii) Ca metal (iii) S (iv) Cl2

(i) Reaction with Ethene CH2 = CH2(s) + H2(g) CH3 – CH3

(ii) Reaction with Ca metal Ca(s) + H2(g) Ca+2 H2(s) –

(iii) Reaction with S H2(g) +S(s) H2S(g)

(iv) Reaction with Cl2 H2(g) +Cl2(g) 2HCl(g)

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Q - 6) Define Water and describe some main physical and chemical properties of Water:

Water: Water is the most abundant compound on Earth's surface, covering

about 70 percent of the planet. In nature, water exists in liquid, solid, and

gaseous states.

Physical Properties of Water: 1. It is a colourless, odourless and tasteless liquid and excellent solvent.

2. It melts and boils at much higher temperature than other liquids.

3. At 4 oC its density is maximum about 1 g/cm3.

4. It freezes at 0 °C and boils at 100 °C.

5. It is one of the few substances that expand upon freezing.

6. Water does not obey law of expansion.

Chemical Properties of Water:

Alkali metals like Na, K etc. and alkaline earth metals like Ca, Ba react with cold water to form their hydroxides with the liberation of H2 gas.

Cold Sodium Hydroxide

1. 2Na(s) + H2O(l) 2NaOH(aq) + H2(g) Cold Calcium Hydroxide

2. Ca(s) + 2H2O(l) Ca(OH)2(aq) + H2(g) (slow) Cold Potassium Hydroxide

3. 2K(s) + H2O(l) 2KOH(aq) + H2(g)

Less electropositive metals like Mg, Zn or Fe etc. reacts with hot water to liberate H2 gas with formation of their oxides. Iron reacts with excess of steam at red heat.

Hot 1. Mg(s) + H2O(l) MgO(s) + H2(g) Hot 2. Zn(s) + H2O(l) ZnO(s) + H2(g) Red Heat 3. 3Fe(s) + 4H2O(g) Fe3O4(s) + 4H2(g)

Noble metals like copper, gold, silver and mercury do not react with water

in any form.

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a) With Chlorine Chlorine reacts with water to produce HCl and hypochlorous acid

(HClO). Hypochlorous acid is unstable and readily liberates atomic oxygen, which can bleach dyes and kill bacteria by oxidation. Chlorine, therefore in water is both bleaching as well as oxidizing agents.

Cl2(s) + H2O(l) ⇌ HCl(aq) + HClO (aq)

HClO(aq) 2HCl(aq) + [ O ] O +O O2(g)

b) With Carbon When steam is passed over heated coke at 1000 oC, a mixture of hydro-

gen and carbon monoxide, known as water gas is produced. 1000 oC

C(s) + H2O(g) CO(g) + H2(g) ⏟

𝑊𝑎𝑡𝑒𝑟 𝐺𝑎𝑠 c) With Silicon

Silicon reacts with steam at very high temperature to form an oxide of silicon i.e. silicon dioxide with liberation of H2 gas.

High temp.

Si (s) + 2H2O(g) SiO2 (g) + 2H2(g) Calcium oxide partially dissolves in water to form calcium hydroxide

i.e. slaked lime. CaO(s) + H2O(l) Ca(OH)2(s) Quick lime Slaked lime

Nitrous oxide dissolves in water to form a neutral solution. N2O(g) + H2O(l) Dissolves to form a neutral solution

(Nitrous oxide)

Q - 7) What do you understand by the anomalous behaviour of water? And what is the significant of this unusual behaviour of water?

Anomalous behaviour of Water: water does not obey the law of expansion and contrac-

tion between 0 oC to 4 oC and shows anomalous or unusual behaviour because water

is polar molecule and due to greater polarity all water molecules are associated by

means of hydrogen bonding. It expands when cooled form 4 to 0 oC and contracts

from 0 to 4 oC. Due to this ice (Solid Water) floats over water and aquatic animals in

winter seasons in the region where temperature reaches below 0 oC.

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Q - 8) What is water of crystallization? What happens when hydrates are heated? Define heat of hydration also write formulas of some of hydrates.

Water of Crystallization: Water of crystallization is water that contains few

molecules of water as a part of the crystal lattice. Crystal salts containing

water of crystallization are called hydrates.

When hydrates are heated: When hydrates are heated, the molecules of water of crystallization are easily dissociated form their salts crystals. The residue left behind is then said to be anhydrous (shapeless) or anhydrite.

CuSO4.5H2O(s) CuSO4(s) + 5H2O(g)

Blue crystals evaporation Anhydrous (white powder)

Heat of Hydration: The minimum amount of heat liberated in the formation of hy-drate is called heat of hydration. Formulas of Hydrates: 1. Copper Sulphate (CuSO4. 5H2O) 2. Ferrous Sulphate (FeSO4. 7H2O)

3. Sodium Carbonate (Na2CO3. 10H2O) 4. Aluminium chloride (AlCl3. 6H2O)

5. Barium Chloride (BaCl2. 2H2O) 6. Potash Alum [K2SO4. Al2(SO4)3. 24H2O]

Q - 9) What is “Potable water”? Write four main characteristics of it.

Potable water: The water that is fit for drinking purpose is called potable water. Characteristics of Portable water:

1. It should be free form germs, bacteria and all sorts of pollutants. 2. It should be moderately soft and its hardness should be under 150 ppm (parts per million) 3. It should be free from corrosive substances. 4. It should be colour less odourless and tasteless and have pH rage of 7 – 8.5. Q - 10) What do you mean by soft water, hard water and heavy water?

Soft water: The water that containing dissolved impurities but in small quantities and easily produces lather with soap is known as soft water.

Hard water: The water that containing dissolved impurities of hydrogen car-bonates, chlorides and Sulphate of calcium and magnesium is called hard water. It does not give lather with soap.

Heavy water: The water that containing compound of oxygen with heavy hydrogen i.e. deuterium is known as soft water. Its formula is D2O.

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Uses of Heavy water:

1. It is used as moderator in nuclear fission reactions to slow down the neutrons. 2. It is used as tracer in biological and chemical researches. Characteristics/Properties of Heavy water:

1. Its density is slightly greater than ordinary water and is 1.104g/cm3. 2. It has low vapour pressure than ordinary water. 3. Its melting point is 3.81 0C and boiling point 101.42 0C. 4. The molecular mass of heavy water is 20 a.m.u. Q - 11) What do you mean by hard water and describes the types of hardness? Also write causes of hardness.

Hard water: The water that containing dissolved impurities of hydrogen car-bonates, chlorides and Sulphate of calcium and magnesium is called hard water. It does not give lather with soap. Types of Hardness:

There are two types of hardness in water. 1. Temporary hardness 2. Permanent hardness Temporary Hardness: Temporary hardness is due to the presence of dissolved hydrogen carbonates of calcium and magnesium. These salts are water soluble and ionize into water as:

Ca(HCO3)2(aq) ⇌ Ca2+(aq) + 2HCO–

3(aq)

Mg(HCO3)2(aq) ⇌ Ma2+(aq) + 2HCO–

3(aq)

Permanent Hardness: Permanent hardness is due to dissolved chlorides and Sulphates of calcium and magnesium. These salts are also water soluble and ionize into water as:

CaCl2(aq) ⇌ Ca2+(aq) + 2Cl–(aq)

Mg(SO4)(aq) ⇌ Ma2+

(aq) + SO42–

3(aq)

Causes of Hardness: 1. Presence of bicarbonates of chlorides or Sulphates of calcium or magnesium

in water is the main causes of hardness of water. 2. Water hardness is caused by the presence of calcium Sulphate (Ca2+) and

magnesium hydroxide (Mg2+) ions, when they occur in very high concentra-tions.

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Q - 12) How is the hardness of water removed? Describe various meth-ods to remove hardness of water?

Method to remove hardness of water: The hardness of water can be removed by the following methods. 1. Temporary hardness

a) By heating b) By Clark’s method

2. Permanent hardness a) By Ion – exchange method

b) By using washing Soda

c) By using caustic Soda

d) By using zeolite or permutit Removing Temporary Hardness

a) By Heating: Temporary hardness is due to the presence of dissolved hydrogen carbonates of calcium and magnesium which decompose on heating to CaCO3 and MgCO3 which are insoluble in water and easily remove by filtration.

Ca(HCO3)2(aq) 𝐵𝑜𝑖𝑙 →

CaCO3 (↓𝑠 )(insoluble) + CO2(g) +

H2O(l)

Mg(HCO3)2(aq) 𝐵𝑜𝑖𝑙 → MgCO3 (

↓𝑠) (insoluble) + CO2(g) + H2O(l)

Once Ca2+ ions or Mg2+ ions are out of water, any soap when added to water becomes soft, the water and lather formation occurs. b) By Clark’s method: Temporary hardness can also be by using Clark’s method. In this method temporary hard water containing hydrogen carbonates of Ca and Mg is treated with slaked lime in the tanks to convert them into their insoluble carbonates. These insoluble carbonates settle down at the bottom of tanks while soft water is drained off for the use.

Ca(HCO3)2(aq) + Ca(OH)2aq 2CaCO3 (insoluble) + 2H2O(l)

Mg(HCO3)2(aq) + Ca(OH)2aq MgCO3 (insoluble) + 2CaCO3 (insoluble) + 2H2O(l)

Removing Permanent Hardness

a) By Ion – exchange method: In this method in which calcium and magnesium ion from water are removed as insoluble precipitates. The chemicals employed are mostly soluble sodium com-pounds. These form insoluble precipitates of Ca and Mg ions. b) By Using washing soda: (Na2CO3 10H2O)

When washing soda is added to permanent hard water, insoluble CaCO3 and MCO3 are precipitated from the soluble salts of Ca and Mg.

CaSO4 (aq) + Na2CO3(aq) CaCO3(s) (insoluble) + Na2SO4(aq)

MgCl2 (aq) + Na2CO3(aq) MaCO3(s) (insoluble) + 2NaCl (aq)

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c) By Using Caustic Soda: (NaOH) When caustic soda is added to permanent hard water, insoluble hydroxide of Mg2+ ion is precipitated from the salts of Ca(OH)2 is partially soluble in water.

MgSO4 (aq) + 2NaOH (aq) Ma(OH)2(s) (insoluble) + Na2SO4(aq)

d) By Using Zeolite or Permutit:

This is a modern method employed for the softening of hard water. Hydrated sodium Aluminium silicate (Na2Al2Si2O8.xH2O) is called permutit. These complex salts are also known as zeolites.

The permutit as loosely packed in a big tank over a layer of coarse sand. Hard water is introduced into the tank from the top. Water reaches the bot-tom of the tank and then slowly rises through the permutit layer in the tank and hardness is removed. The cat ions present in hard water are exchanged for sodium ions. Therefore this method is also called ion exchange method.

Where Z = Al2Si2O8.nH2O CaSO4 (aq) + Na2Z(zeo-

lite) CaZ(s) (insoluble) + Na2SO4(aq)

Sodium zeolite can be regenerated by passing a strong NaCl solution through Ca – Zeolite. CaZ (zeolite) + 2NaCl(aq) Na2Z(s) (Sod – zeolite) + CaCl2(aq)

Q - 13) Describe disadvantages of hard water and uses of water?

Disadvantages of hard water: 1. It consumes more soap in washing process. 2. It may cause stomach disorder or dysentery whenever drinks. 3. If hard water is used in steam engine and turbine then more fuel is used for their

heating. 4. If hard water is used in boiler then it reduces the strength of the boiler metal. Uses of water

1. Water is vital for life. Humans, plants and animals cannot survive without water.

2. It is required for irrigating crops, as seeds cannot germinate without water.

3. Water is used in cooking and washing. Running water is used to generate elec-

tricity.

4. Water serves as a medium for transportation, as ships and boats move on water.

5. Many industries such as petroleum, fertilizer, dye and drugs industries require

large quantities of water for various processes.

6. Fish, other animals and many plants live in water. Fishing and other water-related

activities are a source of livelihood for many people.

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Q - 14) Describe the classification of water Pollutants: or name water borne diseases that are caused by microorganism presents in water, name various types of water pollutants and their different categories.

Classification of water Pollutants: There are various types of water pollutants which can broadly be classified as (a) Oxygen demanding wastes (b) Synthetic organic compounds (c) disease – causing wastes (microorganisms) (d) agricultural water pollutants. a) Oxygen demanding wastes: (OD) These include domestic and animal sewage, bio - degradable organic compounds and industrial wastes from food – processing plants, meat packing plants, slaughter houses, paper and pulp mills, tanneries etc. all these wastes undergo degradation and decompo-sition due to which there is a rapid depletion of demand oxygen that is harm for aquatic

animals. b) Synthetic organic compounds: These are the man – made materials such as synthetic pesticides, synthetic detergents, food additives, pharmaceuticals insecticides, paints, fibres, solvents, plastics etc. these materials are potentially toxic to plants, animals and humans. They cause offensive col-ours, odours and tastes in water. c) Disease causing wastes: (Microorganisms) These include pathogenic microorganisms which may enter water along with sewage and other wastes and may cause tremendous damage to public health. These microbes comprising mainly of viruses and bacteria can cause dangerous water – borne diseases

such as typhoid, cholera, polio, dysentery, and infections hepatitis in humans. d) Agriculture water pollutants: In modern agriculture, pesticides, fertilizers and organic wastes (manure) are essential for producing high yields of crops required for the worlds’ growing population. Some com-

mon pesticides used in Pakistan are alder in, DDT, dielderin etc. Q - 15) Describe hygroscopic substances:

Hygroscopic Substances: The substances which absorb moisture on exposure to atmosphere are called Hygroscopic Substances. CHARACTERISTICS OF HYGROSCOPIC SUBSTANCES In Solid state they do not form solution but merely become sticky or moist when exposed to

atmosphere. Such as sodium nitrate (NaNO3), copper oxide (CuO), Quick lime (CaO).

In Liquid state they absorb water from the atmosphere usually diluting itself up to about three

times of its original volume. Such as Sulphuric Acid (H2SO4).

Uses: They are commonly used in laboratory as drying agents.

15

Q - 16) Name only some common treatment to make municipal water fit for drinking purposes.

The names of some common treatment to make municipal water fit for

drinking are as follows:

(a) Aeration (b) Settling (c) Coagulation (d) Filtration (f) Chlorination Q - 17) What happens when? (i) Mg metal is reacted with hot water.

Ans: When Mg metal is reacted with hot water it produces MgO and lib-

erates H2 gas. Hot

Mg(s) + H2O(l) MgO(s) + H2(g) (ii) Methane is heated above 700 0C in the absence of air.

Ans: When Methane is heated above 700 0C in the absence of air it produces

Carbon black [C(s)] and liberates H2 gas.

Above 700 oC

CH4(g) C(s) + 2H2(g) Absence of air Carbon Black

(iii) Water gas is heated under pressure in the presence of ZnO – Cr2O3 catalyst.

Ans: When water gas (H2) is heated under pressure in the presence of

ZnO – Cr2O3 catalyst it produces methyl alcohol.

ZnO/Cr2O3

CO(g) + 2H2(g) CH3 – OH(l)

Water gas 400oC /high presusure Methyl Alcohol

(iv) A piece of Zn metal is added to the acidic solution of FeCl3.

Ans: When Zn metal is added to the acidic solution of FeCl3 it produces

ferrous chloride (FeCl2) and hydrochloric acid (HCl). Zn

FeCl3(aq) + [H] FeCl2(aq) + HCl(aq)

Nascent Hydrogen

Ferrous Chloride

16

Q - 1) Describe Carbon and its uses. Also write its physical and chemical properties:

Carbon: Carbon is a pure non – metal and the sixteenth most abundant element in the earth crust. Carbon belongs to IV – A group its atomic number is 6 and atomic weight is 12. Carbon exists in three crystalline i.e. graphite, diamond and Bucky balls. In the Free State carbon occurs in following forms:

1. Crystalline forms: a. Diamond b. Graphite c. Bucky balls

2. Amorphous forms: a. Lamp black b. Wood charcoal c. Coal

d. Animal charcoal e. Gas Carbon Physical Properties of Carbon:

1. It is odourless and tasteless solid. 2. All the different carbon allotropes are black or greyish black solids except diamond. 3. It has high melting point above 3000 0C. 4. It is insoluble in all common solvents like water, alcohols, acids and petrol etc. That’s why Carbon deposits formed during incomplete combustion of fuels (Petrol)

inside motor engines have to be removed mechanically; this process is called decarbonizing

of motor engines.

Uses of Carbon: As Diamond: 1. It is used as gems and precious stones because of its sparkling brilliance. 2. Black diamond is used in drillings, in making of instruments for cutting glasses and metals. 3. Its tiny fragments are used as abrasive for polishing tools. As Graphite: 1. It is used as lubricant to reduce friction in machines, bicycle chains and bearings. 2. Its lined crucibles are used for making high grade steel and other alloys. 3. It is also used in lead pencils and as black pigment in paints. 4. It is also used as neutron moderator in nuclear reactions 5. It is used for making inert electrodes in dry cells and in industrial electrolytic

processes.

CHEMISTRY CARBON, SILICON AND THEIR COMPOUNDS

CHAPTER

IX TWELVE

17

As Coal and Coke: 1. Coal and Coke are important fuels and source of energy for homes and industries. 2. It is also used in electric power generator. 3. Coke is used in the extraction of metals from their ores in manufacture of iron and steel. As Charcoal: 1. It is mainly used as a domestic fuel and also as absorbent. 2. Activated charcoal is used in gas masks for absorbing poisonous gases. 3. It is also used for decolorizing of petroleum Gel. As Carbon black: 1. It is used in manufacturing of rubber tyers as filler to increase the strength and hardness of

rubber. 2. It is also used in black shoes polishes, printer’s ink, type – writing papers etc. As Carbon fibre: 1. It is used in manufacturing of incorporated plastics to produce a very light but stiff and strong

material. Chemical Properties of Carbon: 1. Combustion of Carbon (C) All forms of carbon burn in excess of air (O2) to produce carbon dioxide gas (CO2). Its bond disso-ciation energy is – 394 K.J/mol.

C(s) + O2(g) CO2(g) H = – 394 KJ/mol.

In the limited supply of air, incomplete combustion may take place to produce carbon monoxide (CO) instead of CO2 gas.

C(s) + O2(g) 2CO(g)

2. Combination reactions: Carbon combines directly with other elements such as hydrogen, Sulphur, calcium, Aluminium at very high temperature to form addition product.

C (s) + 2H2(g) CH4 (g) (Methane)

C(s) + 2S2(s) CS2(l) (Carbon disulphide)

2C(s) + Ca(s) CaC2(s) (calcium carbide)

3C(s) +4Al(s) Al4C3(s) (Aluminium carbide)

3. As a reducing Agent Carbon also shows greater affinity for oxygen and reduces many metal oxides into free metals.

1. 2CuO(s) + C (s) ℎ𝑖𝑔ℎ 𝑡𝑒𝑚𝑝 → 2Cu(s) +CO2(g) 3. 2ZnO(s) + C(s)

ℎ𝑖𝑔ℎ 𝑡𝑒𝑚𝑝 → 2Zn(s) +CO2(g)

2. Fe2O3(s) + 3C (s) ℎ𝑖𝑔ℎ 𝑡𝑒𝑚𝑝 → 2Fe(s) + 3CO(g) 4. 2PbO(s)+ C(s)

ℎ𝑖𝑔ℎ 𝑡𝑒𝑚𝑝 → 2Pb(s) +CO2(g)

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4. Reaction with strong Oxidizing Agents: Carbon reacts with strong oxidizing agents like hot concentrated nitric acid and cons. Sulphuric acid to liberate CO2 gas.

1. C (s) + 4HNO3 (g) (conc) ℎ𝑜𝑡 → CO2(g) + 4NO2(g) + 2H2O(l)

2. C(s) + 2H2SO4 (g) (conc) ℎ𝑜𝑡 → CO2(g) + 2SO2(g) + 2H2O(l)

Q - 2) Define allotropy and allotropes. Discuss chief allotropic forms of carbon, also write their uses & structure.

Allotropy: The existence of two or more different form of the same element in the same state is called Allotropy.

Allotropes: Allotropes are different forms of the same element in the same state and they have same chemical properties but have different physical properties due to different structures or arrangements of the atom. There are three solid allotropic forms of carbon:

Allotropic forms of Carbon: 1. Diamond 2. Graphite 3. Bucky balls

1. Diamond: It is one of the crystalline forms of carbon. It is found chiefly in

South Africa, Brazil, Australia and India. It may be of blue, green, yellow red or black colour. Black colour diamonds are called Bort or carbonado. STRUCTURE OF DIAMOND: In diamond each carbon atoms is covalently bonded with four other carbon atoms to give basic tetrahedral unit and forming a giant three dimensional molecule. The C – C bond length is 1.54A0 and bond energy is 347 kj/mol. Physical Properties of Diamond:

1. It is transparent and bright in pure state also known as hardest natural substance. 2. Its refractive index is 2.45µ which is very high. 3. It is bad conductor of electricity. 4. It has very high melting point about 3500 0C. 5. Its density is about 3.51 g/cm3. Uses of Diamond: 1. It is used as gems and precious stones because of its sparkling brilliance.

2. Black diamond is used in drillings, in making of instruments for cutting

glasses and metals.

3. Its tiny fragments are used as abrasive for polishing tools.

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2. Graphite: It occurs naturally as Plumbago an opaque black solid. It is found

in Siberia, Canada and Sri Lanka. STRUCTURE OF GRAPHITE: In graphite the carbon atoms form flat layers. Each carbon atom in graphite is linked covalently to three other carbon atoms in the same layer to give basic hexagonal ring. Each layer of carbon atoms viewed as two-dimensional sheet polymer or layer lattice at a distance of 1.42

A0. The distance between the parallel layers is 3.35 0A. Physical Properties of Graphite: 1. It is dark grey colour crystalline solid with dull metallic lustre. 2. It is soft and greasy to feel and leaves black mark on paper. 3. It is less dense than diamond. 4. It has density about 2.2 g/cm3. 5. It is good conductor of electricity. 6. It and has high melting point about 3000 0C. Uses of Graphite: 1. It is used as lubricant to reduce friction in machines, bicycle chains and bear-

ings. 2. Its lined crucibles are used for making high grade steel and other alloys. 3. It is also used in lead pencils and as black pigment in paints. 4. It is also used as neutron moderator in nuclear reactions 5. It is used for making inert electrodes in dry cells and in industrial electrolytic

processes.

3. Bucky Balls: It is a spherical fullerene molecule with the formula C60. It

has a cage-like fused-ring structure (Truncated icosahedron) which resembles a soccer ball, made of twenty hexagons and twelve pentagons, with a carbon atom at each vertex of each polygon and a bond along each polygon edge.

Physical Properties of Bucky Balls:

1. They are resilient to impact and deformation. 2. They are not very reactive due to the stability of the graphite-like bonds. 3. They are also fairly insoluble in many solvents and also conduct electricity. Uses of Bucky balls: 1. The antioxidant properties of Bucky balls may be able to fight the deteriora-

tion of motor function due to multiple sclerosis.

2. Bucky balls may be used to store hydrogen, possibly as a fuel tank for fuel cell powered cars.

3. Bucky balls may be able to reduce the growth of bacteria in pipes and mem-branes in water systems.

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Q - 3) Discus amorphous forms of Carbon. Or Short note

Amorphous For Carbon: The amorphous forms of carbon are not considered as allo-tropes of carbon because x-rays analysis have revealed that they have structures like graphite with the exception of coal which is mined directly from natural depos-its. The other amorphous forms can be prepared in various ways.

Coal: It is originates from the vegetation of the carboniferous era. It is said that

the decomposition of plants and trees occurred gradually under the earth in the absence of air and under pressure CO2 water and methane were liberated leaving behind a material containing high percentage of carbon. During this process under the earth the vegetable material was converted in stages into peat, lignite (Brown Coal), bituminous coal (Soft) and finally to anthracite (Hard Coal). Coal can be mined at various depths from the earth surface and mainly used as fuel.

Coke: it is prefaced by heating bituminous coal to very high temperature about

1300 0C in the absence of air to remove all the volatile constituents present in coal. The process is called destructive distillation of coal. The other non – volatile products obtained by this process would be Coal-Tar and Coke. It burns in air with no smoke and leaves very little residue. Coke is used as fuel and also as reducing agent in the extraction of metals especially iron.

Charcoal: It can be produced by heating wood, nut shells, bones, sugar etc.

Wood charcoal is most common and prepared by burning wood in the limited supply of air. It may contain impurities such as Sulphur. It is mainly used as domestic fuel. Animal charcoal is produced when animal bones and refuse are heated in the lim-ited supply of air it contains high percentage of calcium phosphate Ca3 (PO4)2 as impurity. It is used sugar industries to remove the brown colour from cane sugar and also in decolorizing petroleum Gel. Q - 4) Define catenation:

Catenation: It is the ability of the atoms of carbon to bond itself forming long chains and rings and also to from compounds chain and ring together. This is the unique properties of carbon.

Amorphous Forms of Carbon: (i) Coal (ii) Coke (iii) Charcoal

| | |

— c — c — c —

| | |

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Q - 5) What is silicon? How does silicon occur in nature?

Silicon: It is a metalloid. It belongs to IV-A group in the periodic table. It is the second most abundant element found in the earth’s crust after oxygen. Its atomic number is 14 and atomic weight is 28. OCCURRENCE OF SILICON: 1. Silicon does not occur in Free states, although silicon is widely distributed in

nature. 2. In the combined stated it occurs mainly as silicon IV oxide; SiO2 (Silica) which

is present in various forms. Such as sand, quartz, flint, kieselguhr, agate, etc. 3. Silicon occurs as complex silicates with metallic oxides, like Al2O3, CaO, MgO,

K2O etc.

Q - 6) What are silicates? And Describe some common silicates and their chemical formulas and uses.

Silicates: Silicates are those compounds which have a silicon-oxygen anion chem-

ically combined with such metals as Aluminium, calcium, magnesium, iron, potas-

sium, sodium and others to form silicate salts.

SOME COMMON SILICATES:

Name Chemical Formu-

las Uses

1. Feldspar K2O.Al2O3.6SiO2

or KAlSi3O8

Ceramics, Glass, Pottery and Abrasive

2. Kaolin (China Clay) Al2O3.SiO2.2H2O

Hydrated Crockery

3. Mica

K2O3Al2O3.6SiO2.2H2O

Or KAl3Si2O10

Hydrated

Electrical insulator re-sistance to high temp.

4. Talc (Soapstone) 3MgO.4SiO2.H2O

Hydrated Ceramics

5. Asbestos CaO.3MgO.4SiO2

Or CaMgSi4O12

Heat insulation, Fire – proofing

There are more than 1000 silicates present in the earth’s crust. Kaolin and

china clay consist of hydrated Aluminium silicates.

22

Q - 7) Give some physical properties of silicon. Physical properties of Silicon: 1. Amorphous silicon is brown collared hygroscopic powder, having specific

gravity 2.35.

2. Crystalline silicon is grey in colour, opaque lustrous and octahedral crystal-

line solid with specific gravity 2.49.

3. It is non-volatile solid with very high melting point about 1410 0C & boiling

about 2600 0C.

4. It is hard enough to scratch glass. It is brittle in nature.

5. It is insoluble in most of the common solvents like water but it dissolves in

hydrofluoric acid (HF).

6. It is poor conductor of electricity but sometimes used as semi – conductor.

Q - 8) Describe chemical preparation of Silicon. 1. By heating a mixture of pure dry sand with Mg. Silicon is also prepared by heating a mixture of pure dry sand (SiO2) & magnesium metal in a fire clay crucible in the absence of air.

SiO2(s) + 2Mg(s) ℎ𝑒𝑎𝑡→ Si(s) + 2MgO(s)

Dilute hydrochloric acid is then added in the reaction mixture to dissolve unreacted Mg metal and MgO formed. The residue left behind contains amorphous silicon.

Mg(s) + 2HCl(aq) ℎ𝑒𝑎𝑡 → MgCl2(aq) + H2(g)

MgO(s) + 2HCl(aq) ℎ𝑒𝑎𝑡 → MgCl2(aq) + H2O(l)

If unreacted SiO2 is left, it can also be removed by dissolving it in hydro-fluoric acid (HF). SiO2(s) + 4HF ℎ𝑒𝑎𝑡

→ SiF4(aq) + 2H2O

(Silicon tetra fluoride)

2. By passing vapours of SICl4 over heated sodium or potassium metal. When the vapours of SiCl4 are passed over heated sodium or potassium metal in an inert atmos-phere, silicon is produced by the reduction process.

1. SiCl4(g) + 4Na(s) ℎ𝑒𝑎𝑡 → Si(s) + 2NaCl(s)

2. SiCl4(g) + 4K(s) ℎ𝑒𝑎𝑡 → Si(s) + 4KCl(s)

3. By heating SiO2 with coke Silicon is also prepared by heating SiO2 with coke in an electric furnace. This is an industrial method.

ℎ𝑒𝑎𝑡 →

SiCl4(g) + 2C(s) Si(s) + 2CO(g) Eclectic furnace Crystals

23

Q - 9) Give some uses of silicon: Uses of Silicon: 1. Silicon is used in bronze and steel alloys to increase their tensile strength. 2. Very pure silicon is used in making semi – conductors which are of great

importance in computers, transistors etc. 3. It is also used for making silicones which are rubber like liquids. 4. It is also used as lubricant, water – repellent, electric insulators. 5. It is also used in paints, varnishes and polishes. 6. Silicone is also used in the preparation of refractory material such as cruci-

ble, fire – bricks etc.

Q - 10) Describe preparation of Silica and its properties and uses:

Silica: Silica occurs naturally in three main crystalline forms namely quartz, tridymite and crysto balite. The commonest of three is quartz. Physical properties of Silica: 1. In pure state Silica exists in a colourless crystalline form. 2. It is macromolecular compound with silicon and oxygen atoms. 3. It is non – volatile and hard. 4. Its melting point is about 15000C.

Preparation of Silica: By heating silicon in Air/Oxygen Silica is prepared by heating silicon in air or oxygen.

Si (s) + O2(g) ℎ𝑒𝑎𝑡 → SiO2(s)

It can also be prepared in hydrated form as a gelatinous precipitate by warming sodium sili-

cate (Na2SiO3) with conc. HCl solution.

Na2SiO3(aq) + 2HCl (conc) ℎ𝑒𝑎𝑡 → SiO2.H2O + 2NaCl(aq)

Hydrated Silica

Uses of Silica: 1. It is widely used in making mortar, cement, concrete, glass and refractory silica

bricks. 2. Fused silica (quartz glass) is used in making optical lenses and prisms, heat –

resisting articles etc. 3. Large quartz crystals are used for lenses of optical instruments. 4. Powered quartz is used in the making of silicon carbide (SiC), Silicon tetra flu-

oride (SiF4), sodium silicate (Na2SiO3) and silica bricks for lining furnaces. 5. Kieselguhr (SiO2) absorbs liquids readily and is used as absorbent of nitro-glyc-

erine (explosive) in the making of dynamite. 6. It is also used in medicines for making dry antiseptic dressings.

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Q - 11) Describe preparation of Sodium Silicate or water glass and its properties and uses:

Sodium Silicate or Water Glass: Sodium silicate is the common name for a com-pound sodium Meta silicate, Na2SiO3, also known as water glass.

Or Sodium silicate dissolves in hot water under pressure to form a viscous liquid which is known as water glass because it looks like ordinary glass. Physical properties of Sodium Silica: 1. It is colourless glass-like solid. 2. Its melting point is about 10900C.

3. Its density is 2.40 g/cm3. 4. It is very soluble in water

Preparation of Sodium Silica: By heating silica and sodium carbonate Sodium silicate is prepared by heating strongly two parts by mass of silica i.e. sand (SiO2) with one part by mass of sodium carbonate (Na2CO3) until the mixture melt.

Na2CO3(s) + SiO2(s) 𝑠𝑡𝑟𝑜𝑛𝑔 ℎ𝑒𝑎𝑡𝑖𝑛𝑔 → Na2SiO3(s) + CO2

↑(g)

It is obtained glass-like solid with melting point 10900C.

Uses of water glass: 1. It is used to get silica Gel. 2. It is used as filler in soap industries. 3. It is used for sizing of paper, for fire-proofing of wood and textiles and for making

glue. 4. It is also used for making chemical garden. Q - 12) What is Silica Gel? Also write some uses of it.

Silica Gel: When an acid is added in a solution of water glass. It turns into a Gel like substance known as Gel. The formula for Gel is SiO2.nH2O. On complete dehy-dration by heating a hard porous material is obtained known as Silica Gel.

Or

It is a granular, vitreous, porous form of silicon dioxide (SiO2) made synthetically from sodium silicate (water glass). Silica gel is hard and more solid than common household gels like gelatine or agar.

Uses of Silica gel: 1. It is used as good absorbent and in the refining of petroleum. 2. It is used to prevent medicines being spoiled. 3. It is also used to recover valuable vapours form industrial effluents.

25

Q - 1) Describe Nitrogen, its occurrence and uses. Also write physical and chemical properties:

NITROGEN: Nitrogen is the most common gas present in the atmosphere and the tenth most abundant element in the earth crust. Nitrogen belongs to V – A group its atomic number is 7 and atomic weight is 14. Nitrogen was discovered by Scottish Botanist, Daniel Rutherford in 1772. It does not sustain combustion. OCCURRENCE OF NITROGEN: 1. Nitrogen occurs in the Free State as N2 gas in air up to 78% by volume and 75% by mass. 2. In combine state nitrogen occurs abundantly in the earth’s crust as nitrates of sodium, calcium

and potassium. 3. In combine state nitrogen is found in organic matter such as proteins, urea and vitamin B com-

pounds. In Free State nitrogen in air is important because it dilutes oxygen.

Physical Properties of Nitrogen:

1. It is a colourless, odourless and tasteless gas. 2. Pure nitrogen is slightly soluble in water. 3. It is slightly lighter than air. 4. Its boiling point is -196 0C, while melting points – 210 0C. Uses of Nitrogen: 1. In the form of nitrous oxide it is used as an anaesthetic. 2. Cryopreservation also uses Nitrogen to conserve egg, blood, sperm and other biological

specimens. 3. The CPUs in computers use Nitrogen gas to keep them from heating up. 4. It also serves as an oxidation reaction catalyst. 5. Apart from being an oxidizing agent, it can also be used as a flour bleaching agent and rocket

fuel.

Chemical Properties of Nitrogen: Molecular nitrogen is unreactive because of strong triple bond between two nitrogen at-oms(N ≡ N). Its dissociation energy is 941 Kj/mol. 1. Reaction with Hydrogen (H) At very high temperatures & pressures, nitrogen combines directly with H forms ammonia (NH3)

N2(s) + 3H2(g) 450 0C /200−250 atm → NH3(g)

2. Reaction with Oxygen (O) At 2000 0C temperatures, nitrogen combines directly with oxygen and forms nitric oxide NO(g).

N2(s) + O2(g) 2NO(g)

3. As a reducing Agent At very high temperatures, nitrogen combines directly with magnesium and forms magnesium ni-tride Mg3N2 (g).

N2(s) + 3Mg2(s) Mg3N2(g)

CHEMISTRY NITROGEN AND OXYGEN CHAPTER

IX THIRTEEN

26

Q - 2) Describe Oxygen and its occurrence & uses. Also write its physi-cal and chemical properties:

OXYGEN: Oxygen is the most abundant element in the earth. Oxygen belongs to V–A group its atomic number is 6 and atomic weight is 12. Oxygen was discovered by Scheele in 1772 and Priestley in 1774. However the major properties of oxygen is given by Lavoisier. It is most essential substance for all living things. OCCURRENCE OF OXYGEN: 1. Oxygen occurs in the Free State as well as combined state.

2. In Free State oxygen is present as diatomic gas (O2) in the earth’s atmospheric

air up to 21% by volume and about 33% by mass.

3. In combine state oxygen accounts for nearly 50% by mass of the earth’s crust.

Physical Properties of Oxygen:

1. It is a colourless, odourless and tasteless gas. 2. It is neutral to moist litmus paper. 3. It is slightly soluble in water only about 2% by volume at room temperature. 4. Gaseous oxygen is about 1.1 times denser than air. 5. It liquefies at -183 0C and solidifies at – 210 0C. Uses of Oxygen: 1. This gas is used in various industrial chemical applications. 2. It is used to make acids, sulphuric acid, nitric acid and other compounds. 3. Hot oxygen air is required to make steel and iron in blast furnaces. 4. Some mining companies use it to destroy rocks. 5. Oxygen gas is used to destroy bacteria. Chemical Properties of Oxygen: Oxygen reacts with metal, non-metals and other compounds directly. 1. Reaction with Calcium (Ca)

Oxygen combines directly with Calcium forms Calcium oxides (CaO) 2Ca2(s) + O2(g)

heat→ 2CaO(s)

2. Reaction with Lithium (Li) Oxygen combines directly with Lithium forms Lithium oxides (Li2O)

4Li(s) + O2(g) heat→ 2Li2O(s)

3. Reaction with Ferric sulphide (FeS) Oxygen combines directly with Ferric sulphide forms Ferric oxides (Fe2O3) & Sulphur dioxide.

4FeS2(s) + 7O2(g) heat→ 2Fe2O3(s) + 4SO2(g)

27

4. Reaction with Sulphur (S) Oxygen combines directly with Sulphur forms Sulphur dioxides (SO2)

S(s) + O2(g) iginitio𝑛 → SO2(g)

5. Reaction with Carbon (C) Oxygen combines directly with Carbon forms Carbon dioxides (CO2)

C2(s) + O2(g) iginitio𝑛 → CO2(g)

6. Reaction with methane (CH4) Oxygen combines directly with Methane forms Carbon dioxides and water.

CH4(g) + 2O2(g) combussion → CO2(g) + 2H2O(g)

7. Reaction with Hydrogen sulphide (H2S) Oxygen combines directly with Hydrogen sulphide forms Sulphur dioxides and water.

2H2S(g) + 3O2(g) combussion → 2SO2(g) + 2H2O(g)

Q - 3) Describe chemical preparation of Nitrogen. How can you get ni-trogen from the atmospheric air? Give tow uses of nitrogen. 1. From Air The only important method of producing nitrogen gas is the fractional distillation of liquid air. In this process air is firs liquefied to form liquid air which is ten fractionally distilled. Air liquefied by successive compression and expansion. Fractional distillation of liquid air

1. Clean air is compressed and then cooled by refrigeration, upon expanding the air further cools and liquefies.

2. The liquid air is filtered to remove carbon dioxide solid and then distilled. 3. Nitrogen is the most volatile component, with boiling point -196 oC, distil over. 2. In laboratory Pure nitrogen in the laboratory is prepared by heating ammonium nitrate which thermally decompose to give nitrogen gas. Ammonium nitrite is first ob-tained by reacting ammonium chloride with sodium nitrite.

a. Formation of Ammonium Nitrite NH4Cl(s) + NaNO2(s)

heat → NH4NO2(s) + NaCl(s)

b. Preparation of N2 NH4NO2(s)

heat → N2(g) + 2H2O(l)

Uses of Nitrogen: 1. In the form of nitrous oxide it is used as an anaesthetic. 2. Cryopreservation also uses Nitrogen to conserve egg, blood, sperm and

other biological specimens.

28

Q - 4) Describe laboratory preparation of Oxygen. How oxygen is in-dustrially produced from liquid air? 1. From Air

The isolation of oxygen from air involves two steps. (a) Liquefaction of air (b) Fractional distillation of liquid air.

a. Liquefaction of air

1. Air in the gaseous form is first passed through caustic soda to remove CO2 present in air. 2. It is then compressed under very high pressure about 200 atm in the compressor. 3. It is then cooled and allowed to expand rapidly through a nozzle. 4. The process of compression and expansion are repeated over and over again due to which

temperature falls up to – 200 0C at which air liquefies. b. Fractional distillation of liquid air 1. The liquid air is then led to a fractionating column through a filter in order to remove

the traces of CO2 solid if left behind. 2. On distillation nitrogen with lower boiling point of -190 0C, evolves first leaving behind

a liquid very rich in oxygen. 3. On heating turns liquid argon into gas which boils out at – 185.7 0C and passes off from

the middle of the column and liquid oxygen. 4. The least volatile component in the air turns into oxygen gas at -183 0C.

2. In laboratory Oxygen in the laboratory is prepared by heating potassium chlorate mixed with little manganese dioxide, which acts as a catalyst, the decomposition reaction takes place at lower temperature and at much faster rate.

2KClO3(s) MnO2|heat → 2KCl(s) + 3O2(g)

Q - 5) What are oxides how are they classified describe normal oxides in detail?

OXIDES: The binary compounds of oxygen with metals and non-metals are called as

oxides. Such as CaO, Fe2O3, CO2, H2O etc. Classification of oxides On the basis of valence number or oxidation of oxygen, oxides are classified into several groups.

(a) Normal Oxides (b) Peroxides, (c) Super oxides (d) Sub oxides

a. Normal oxides Normal oxides are those oxides in which oxygen shows normal oxidation state

or valence number -2. It is further divided into four types.

(i) Basic oxides (ii) Acidic oxides (iii) Amphoteric oxides (iv) Neutral oxides

29

(i) Basic oxides The normal oxides of metals are the examples of basic oxides.

1. 2Ca2(s) + O2(g) → 2CaO(s)

2. 2Pb(s) + O2(g) → 2PbO(s)

3. 4Na(s) + O2(g) → 2Na2O(s)

Most of these are soluble in water and produce their hydroxides. They also turns red litmus to blue.

1. CaO(s) + H2O(l) → Ca(OH)2(aq)

2. Na2O(s) + H2O(l) → 2NaOH(aq)

They also react with acids to form salts and water.

3. MgO(s) + 2HCl(aq) → MgCl2(aq) + H2O(l)

4. CaO(s) + 2HNO3(aq) → Ca(NO3)2(aq) + H2O(l)

(ii) Acidic oxides The normal oxides of non-metals are generally acidic.

1. S(s) + O2(g) → SO2(g)

2. C(s) + O2(g) → CO2(g)

3. N2(s) + 2O2(g) → 2NO2(g)

These oxides react with water to form acids which turns blue litmus to red.

5. SO2(g) + H2O(l) → H2SO3(aq) (Sulphurous acid)

6. Na2O5(s) + H2O(l) → 2HNO3(aq) (Nitric acid)

7. CO2(s) + H2O(l) → H2CO3(aq) (Carbonic acid)

They react with alkalis to form salts and water.

8. CO2(g) + 2NaOH(aq) → Na2CO3(aq) + H2O(l)

9. SO3(g) + 2KOH(aq) → K2SO4(aq) + H2O(l)

(iii) Amphoteric Oxides The oxides that possess dual characteristics i.e acidic as well as basics are known as amphoteric oxides.

1. 4Al(s) + 3O2(g) → 2Al2SO3(s)

2. 2Zn(s) + O2(g) → 2ZnO (s)

They react with alkalis (acids/bases) to form salts and water. With acids Al2O3(s) +6HCl(aq)

→ 2Al2Cl3(aq) + 3H2O(l)

ZnO(s) + H2SO4(aq) → ZnSO4(s) + H2O(l)

With bases Al2O3(s) + 2NaOH(aq) → 2NaAlO2(aq) (sod aluminate) + H2O(l)

ZnO(s) + 2NaOH(aq) → NaZnO2(s) (sod zincate)+ H2O(l)

(iv) Neutral Oxides The oxides that are neither acidic nor basic are known as neutral oxides. They are neutral to litmus in aqueous solutions. Example:

1. Nitric oxide (NO)

2. Carbon monoxide (CO)

3. Water (H2O)

4. Nitrous Oxide (N2O)

30

b. Peroxides Peroxides are oxides containing higher proportion of oxygen as compared to normal oxides.

In these oxides Oxygen has oxidation state or valence number 1.

They contain peroxide ion (O – O)

They produce hydrogen peroxide with acids.

Example: 1. Sodium peroxide (Na2O2) 2. Barium Peroxide (BaO2) Na2O2(s) + 2HCl(aq) NaCl(aq) + H2O2(aq) (Hydrogen peroxide)

c. Super oxides Super oxides are oxides containing higher proportion of oxygen as compared to peroxides.

In these oxides

Oxygen has oxidation state or valence number −0.5 or −1

2 .

They show tendency to release oxygen (O2) on heating and powerful oxidizing agent.

They do not produce hydrogen peroxide with acids. Example: 1. Potassium superoxide (KO2)

2. Caesium superoxide (CsO2)

3. Rubidium superoxide (RbO2) d. Sub oxides Sub oxides are oxides containing less quantity of oxygen than normal oxides. In these oxides

They are unstable.

Very few sub oxides are known. Example: Carbon sub oxide (C3O2)

Q - 6) Give the preparation and properties of hydrogen peroxide and its uses:

Hydrogen peroxide (H2O2): The Nard was the first to prepare hydrogen peroxide by the

action of dilute Sulphuric acid on barium peroxide (BaO2). He discovered that hydrogen peroxide contain one more oxygen atom in its molecule than water thus called it as

oxygenated water. H2SO4(aq) + BaO2(s) Heat → BaSO4(s) + H2O2(aq)

Physical Properties of Hydrogen Peroxide:

1. Pure Hydrogen peroxide is a pale blue syrupy liquid. 2. It mixes with water to give solution which is slightly acidic. 3. Its boiling point is 150 oC but it boils with decomposition. 4. It freezes at about – 0.9 oC.

31

Preparation of Hydrogen Peroxide: Hydrogen peroxide is prepared by two Methods.

(i) Laboratory Method (ii) Industrial Method

(i) Laboratory Method: The hydrogen peroxide is usually prepared in laboratory by the action of dilute Sulphuric acid on peroxides of certain metals, especially barium peroxide (BaO2). Barium peroxide is insoluble and can be easily removed by filtration and pure H2O2 is obtain.

H2SO4(aq) + BaO2(s) Heat → BaSO4(s) + H2O2(aq) ← (𝐻𝑦𝑑𝑟𝑜𝑔𝑒𝑛 𝑃𝑒𝑟𝑜𝑥𝑖𝑑𝑒)

(ii) In Industrial Method/Preparation: The hydrogen peroxide is usually prepared in laboratory by the action of dilute Sulphuric acid on peroxides of certain metals, especially barium peroxide (BaO2). Barium peroxide is insoluble and can be easily removed by filtration and pure H2O2 is obtain.

CH3 −CH|OH

− CH3(l) + O2(g) → H2O2(l) + CH3 −

O||C− CH3(l)

Chemical Properties of H2O2: 1. As showing exothermic reaction

When hydrogen peroxide is exposed to air, it decomposes to form water and oxygen. The de-composition is exothermic.

2H2O2(l) exposure → 2H2O(l) + O2(g) + ∆H

2. As Oxidizing agent Hydrogen peroxide is common oxidizing agent, usually in the form of aqueous solution with 3% H2O2. It is strong oxidizing agent because it can readily donate oxygen or accept electrons.

H2O2(l) → H2O(l) + O(g) (Donation of Oxygen)

H2O2(l) + 2H++ 2e– → H2O(l) (Acceptor of Electron)

3. As reducing agent Hydrogen peroxide in the form of aqueous solution with 3% H2O2 can also behave as a reducing agent. It reduces acidic potassium permanganate solution recoloring of KMnO4.

2KMnO4(aq) 3H2SO4(aq)+ 5H2O2(l) → K2SO4(aq)+ 2MnSO4(aq) +H2O(l) + 5O2(g)

It reduces chlorine to hydrochloric acid and Oxygen gas is given off.

H2O2(aq) + Cl2 → 2HCl + O2(l)

Isopropyl alcohol Hydrogen peroxide Acetone

32

Uses of Hydrogen peroxide (H2O2):

1. It is used as a mild antiseptic in mouth wash as well as for cleaning wounds. 2. It is used as bleaching agent in bleaching delicate materials like silk, wool, feathers and

human hairs. 3. It removes unwanted colour from fabrics, hair or other materials. 4. Its liquid is used for restoring painting & providing oxygen for burning fuel in space

rocket. 5. It is used in the preparation of compounds like sodium chlorate (NaClO2).

Q - 7) Define oxidation and reduction and write chemical equation in support of each.

Oxidation: The process or reaction in which oxygen combines with other elements or compounds to produce oxides. It is also known as addition of oxygen.

Reduction: The process or reaction in which removal of oxygen occurs from a substance. It is also known as removal of oxygen. Oxidation involves

(a) Addition of Oxygen (b) Removal of Hydrogen (c) Loss of electrons

a. As addition of oxygen

When oxygen reacts with iron, magnesium or carbon produces their oxides. 4Fe(s) + 3O2(g)

→ 2Fe2O3(s) (Ferric oxide)

Mg(s) + O2(g) → 2MgO (s) (Magnesium oxide)

C(s) + O2(g) → CO2(g) (carbon dioxide)

b. As removal of hydrogen

By this process removal of hydrogen occurs form a compound. H2S(g) + Cl2(g)

→ S(s) + 2HCl

c. As loss of electron

By this process loss of electron occurs from a substance. Sn(s)

→ Sn2+

+ 2e –

Reduction involves (a) Addition of Hydrogen (b) Removal of Oxygen (c) Gain of electrons

a. As addition of Hydrogen

By this process Hydrogen is added to a substance. H2S(g) + Cl2(g)

→ S(s) + 2HCl

b. As removal of Oxygen

By this process removal of oxygen occurs form a compound. CuO(s) + H2(g)

→ Cu(s) + H2O(l)

c. As gain of electron

By this process gain of electron occurs from a substance.

Al3+ + 3e–

→ Al(s)

33

Q - 8) Define oxidizing agent, reducing agent and redox reactions.

Oxidizing Agent: A substance that accepts or gains electrons is known as oxidizing agent. And a substance itself reduced.

Reducing Agent: A substance that donates or loss electrons is known as reducing agent. And a substance itself oxidized.

Redox Reaction: The reaction in which oxidation & reduction occur simultaneously is called oxidation-reduction reaction or redox reactions. Q - 9) What is ozone? How ozone is produced in the atmosphere? Or how oxygen is converted into ozone? Its physical/chemical properties and what is important of ozone and write uses of ozone.

OZONE: Ozone is pale-blue poisonous gas with a sharp, irritating odour. It is an allotropic form of oxygen with molecular formula O3. It was first discovered by Schonbein in 1839. OCCURRENCE OF OZONE: 1. It exists in a layer at a height of about 20 kilometres above the earth. 2. Very small amount of ozone is produced around electrical machineries when

they are in operations.

Physical Properties of Ozone: 1. It is a bluish, collared gas that has a boiling point of -112 oC. 2. It is very poisonous gas at concentration 100 parts per million (ppm). 3. It is only slightly soluble in water but dissolves in turpentine oil readily. 4. It has characteristic smell which is sharp irritating like Cl2 gas. Uses of Ozone: 1. It is used in treatment of domestic water in place of chlorine. 2. It is used as bleaching because all oxidizing agents are also good bleaching agent. 3. It is largely used in the preparation of pharmaceuticals, synthetic lubricants and other

commercially useful organic compounds. Important of Ozone:

It protects the earth from the harmful effects of high energy rays. But In lower atmosphere ozone is measured as air pollutant it damages living system. odour

Chemical Properties of Ozone: Ozone is chemically more reactive than ordinary diatomic oxygen. It acts as pow-erful oxidizing agent because Ozone dissociates readily forming reactive oxygen atoms.

O3 → O2 + O(g) ∆H = −107 KJ/Mol.

34

1. Reaction with Lead sulphide (PbS) Ozone oxidizes lead sulphide (PbS) in acidic medium liberating oxygen (O2) gas.

PbS2(s) + 4O3(g) → PbSO4(s) + 4O2(g)

2. Reaction with Sulphuric Acid (H2SO4) Ozone oxidizes Hydrogen sulphide (H2SO4) in acidic medium liberating oxygen (O2) gas.

H2SO4(s) + 2O3(g) → H2SO4(aq) + O2(g)

3. Reaction with Sulphur dioxide (SO2) Ozone oxidizes Sulphur dioxide (SO2) in acidic medium liberating oxygen (O2) gas.

SO2(s) + O3(g) → SO3(g) + O2(g)

4. Reaction with Potassium Iodide (KI) Ozone oxidizes Potassium Iodide (KI) in acidic medium liberating oxygen (O2) gas.

KI(aq) + O3(g) + H2SO4(s) → K2SO4(aq) + I2(aq) + H2O(l)+ O2(g)

Preparation of Ozone (O3): Ozone is prepared by two Methods.

(i) Atmospheric Method (ii) By electric discharge Method

(i) Atmospheric Method: In nature ozone is formed form atmospheric oxygen by lightning flashes however ozone is very unstable dissociates readily forming reactive.

O3 → O2 + O(g) ∆H = −107 KJ/Mol.

(ii) By Electric discharge method: Ozone can be prepared from oxygen by passing electric discharge through oxygen gas. It is neces-sary to use silent discharge because sparking would generate heat energy which decomposes ozone proceed. The apparatus used for converting oxygen into ozone is known as Ozonizer.

3O2(g) Electric discharge → 2O3(g)

Q - 10) What is Aqua Regia? How does it dissolve gold?

Aqua Regia: The mixture of concentrated nitric acid (HNO3) and hydrochloric acid HCl, optimally in a volume ratio of 1:3 is called Aqua Regia which is also known as "Royal Water". Aqua Regia dissolves gold due to liberation of nascent chlorine which forms gold chloride with it, which is soluble. HNO3(conc) + 3HCl(conc)

→ NOCl + 2H2O(l) + 2Cl

NOCl → NO + Cl

Au(s) + 3Cl → AuCl3

Oxy-

gen Ozon

e

Nitrosyl chlo-

ride Nascent chlo-

rine

Nitrosyl chlo-ride

35

Q - 11) Give the preparation and properties of Ammonia (NH3) and its uses:

Ammonia (NH3): Ammonia is a very important chemical in industry, in nature ammonia

is produced during the decay of nitrogenous matter in the absence of air. Physical Properties of Ammonia: 1. It is colourless gas with a characteristic pungent smell. 2. It is highly soluble in water about 1300ml dissolves in 1ml of water. 3. It is easily liquefied into a colourless liquid at ordinary temperature by com-

pression.

4. In large quantity it is poisonous because of its effects on respiratory system.

Preparation of Ammonia: Ammonia is prepared by two Methods.


(i) Laboratory Method: In the laboratory ammonia is prepared by heating ammonium salts usually ammonium chloride (NH4Cl) with slaked lime i.e. calcium hydroxide.

2NH4Cl(s) + Ca(OH)2(s) Heat → CaCl2(g) + 2H2O(l) + 2NH3(g) ← (Ammonia)

(ii) In Industrial Method/Preparation: (Haber – Bosch process) On large scale ammonia is manufactured by the direct combination of Nitrogen and Hydrogen from Haber – Bosch process. In this process a mixture of pure Ni-trogen and Hydrogen in the ratio of 1:3 by volumes is allowed to react. The basic problem in ammonia syn-thesis is that it is a reversible reaction and can be described as

To get maximum yield of ammonia

The Optimum temperature should be 400 – 450 oC.

The pressure should be 200 – 250 atm.

The suitable catalyst Fe2O3 (Ferric oxide) with small amount of Al2O3, CaO K2O. Uses of Ammonia: 1. Aqueous Ammonia is used in softening of temporary hard water.

2. It is used as solvent in laundries for removing grease & oil stains.

3. It is used as cooling agent in some refrigerators.

4. It is used in manufacturing of nitrogenous fertilizers, like Urea, (NH4)2SO4 etc.

36

Chemical Properties of Ammonia: 1. Reaction with Water:

Ammonia reacts with water to form ammonium hydroxide. NH3 + H2O

→ NH4HO(aq)

2. Reaction with Oxygen: Ammonia reacts with oxygen (O) to form Nitrogen gas (N) and wa-ter vapours.

4NH3 + 3O2 ignition → 2N2(g) + 6H2O(g)

3. Reaction with Acids: Ammonia reacts with acids to form ammonium salts.

a. 2NH3 + H2SO4(aq) → (NH4)2SO4(aq)

b. NH3 + HCl(aq) → NH4Cl(𝑎𝑞)

c. NH3 + HNO3(aq) → NH4NO3(𝑎𝑞)

4. Reaction with Chlorine: Ammonia reacts with chlorine, it first reduces chlorine to produce hydrogen chloride and nitrogen then hydrogen chloride reacts with excess of ammonia to form dense white fumes of ammonium chlo-ride.

a. 2NH3 + Cl2(g) → N2(g) + 6HCl(g)

b. 6NH3(g) + 6HCl(g) → 6NH4Cl(𝑠)𝑤ℎ𝑖𝑡𝑒 𝑓𝑢𝑚𝑒𝑠

The overall reaction can be described as: c. 8NH3(g) + 3Cl2(g)

→ 6NH4Cl(𝑠) + N2(g)

If chlorine is in excess then nitrogen tri-chloride (NCl3) would produce which is an oily liquid and is dangerously explosive. NH3(g) + 3Cl2(g)

→ NCl3(𝑙) + 3HCl(g)

5. Reaction with Carbon dioxide: Ammonia reacts with CO2 at high temperature about 150oC under pressure to produce Urea (NH3)2CO, which is an important fertilizer.

2NH3 + CO2(g) 150°C → (NH2)2CO(s) + H2O(l)

6. As a reducing agent: Ammonia is not a strong reducing agent. However it reduces heated copper oxide to free copper metal with evolution of N2 gas and water.

3CuO(s) + 2NH3(g) heat → 3Cu(s) + N2(g) + 3H2O(l)

High pressure

Urea ↓

37

Q - 12) Give the preparation, properties of Nitric Acid and its uses:

Nitric Acid (HNO3): Nitric acid is a very important acid which is used extensively in the

laboratories and in industries. It was first prepared by Glibber in 1685 from sulphuric acid and potassium nitrate.

Physical Properties of Nitric Acid: 1. It is colourless fuming liquid with sharp chocking smell. 2. Its density is 1.52 g/cm3. And its specific gravity is 1.41.

3. It is miscible in water in all proportion and it has sour taste. 4. The boiling point of pure Nitric Acid is 83oC, freezing point is – 41.6oC.

Preparation of Nitric Acid:

Nitric Acid is prepared by two Methods. (i) Laboratory Method (ii) Industrial Method

(i) Laboratory Method: In the laboratory Nitric Acid is prepared by heating solid Potassium Nitrate (KNO3) with conc. Sulphuric acid (H2SO4).

KNO3 + H2SO4(aq)(conc) Heat → KHSO4(aq) + 2NH3(g) ← ( Nitric Acid)

(ii) In Industrial Method/Preparation: (Ostwald’s Method) On large scale Nitric Acid is manufactured by the catalytic oxidation of ammonia in Ostwald’s Method. In this method ammonia is allowed to react with excess of air in the presence of platinum, catalyst at 600oC it produce Nitric Oxide.

4NH3 + 5O2 pt−600°C → 4NO(g) + 6H2O(l) ; ∆H = −95.5 Kj/mol

Nitric acid is then mixed with excess of air in oxidation chamber to produce nitrogen dioxide i.e. Nitrogen peroxide (NO2). 2NO2(g) + O2(g)

→ 2NO2(g)

Nitrogen peroxide produced is directly dissolve in water to produce nitric acid, liberating nitric oxide gas in the absorption chamber. It recycled to get more nitric acid. 3NO2(g) + H2O(l)

→ 2HNO3(aq) + NO(g)

Uses of Nitric Acid: 1. It is used as laboratory reagent.

2. It is used in the manufacture of cellulose lacquers and smokeless gun powder.

3. It is used in the manufacture of dyes and explosives such as nitro glycerol and

trinitrotoluene (T.N.T) which is powerful explosive.

4. It is used in manufacturing of fertilizers, like NH4NO3, NaNO3, and KNO3 etc.

Potassium Hydro-gen Sulphate

38

Chemical Properties of Nitric Acid: 1. Reaction with Water:

Nitric Acid is a strong monobasic acid reacts with water to form hydronium ion. HNO3 + H2O ⇌ H3O

+ + NO3−

2. Reaction with Metal Oxides/Alkalis/Metal Carbonates: Nitric Acid reacts with metal oxides, alkalis and metal carbonates to form nitrate salts and water. 1. HNO3(aq) + NaOH(aq)

→ NaNO3(aq) + H2O(l)

2. HNO3(aq) + KOH(aq) → KNO3(aq) + H2O(l)

3. CaO(s) + 2HNO3(aq) → Ca(NO3)2(aq) + H2O(l)

4. PbO(s) + 2HNO3(aq) → Pb(NO3)2(aq) + H2O(l)

5. CaCO3(s) + 2HNO3(aq) → Ca(NO3)2(aq) + CO2(g) + H2O(l)

6. NaHCO3(s) + HNO3(aq) → NaNO3(aq) + CO2(g) + H2O(l)

Nitric Acid reacts with some metals like Mg and Mn to form nitrate salts and liberate H2 gas. a. Mg(s) + 2HNO3(aq)

→ Mg(NO3)2(aq) + H2(g)

b. Mn(s) + 2HNO3(aq) → Mn(NO3)2(aq) + H2(g)

3. Reaction with Non – metal: Hot conc Nitric Acid reacts with many non-metals and oxidizes them into their oxides or oxy-acids and itself is reduced to NO2 gas.

a. C(s) + 4HNO3(aq) hot → CO2(g) + 4NO2(g) + 2H2O(l)

b. Si(s) + 4HNO3(aq) hot → SiO2(s) + 4NO2(g) + 2H2O(l)

c. S(s) + 4HNO3(aq) hot → SO2(g) + 4NO2(g) + 2H2O(l)

d. S(s) + 6HNO3(aq) hot → H2SO4(aq) + 6NO2(g) + 2H2O(l)

e. P(s) + 5HNO3(aq) hot → H3PO4(aq) + 5NO2(g) + H2O(l)

f. I2(s) + 10HNO3(aq) hot → 2HIO3(aq) + 10NO2(g) + 4H2O(l)

4. Reaction with Metals: Conc as well as dil. Nitric Acid oxidizes many metals. Conc HNO3 liberates NO2 gas, while the reaction of dil. HNO3 depends upon the nature or the reducing agent. With Copper:

1. Cu(s) + 4HNO3(conc) hot → Cu(NO3)2(aq) + 2NO2(g) + 2H2O(l)

2. 3Cu(s) + 8HNO3(dil.) hot → 3Cu(NO3)2(aq) + 2NO(g) + 4H2O(l)

39

With Lead: 1. Pb(s) + 4HNO3(conc)

hot → Pb(NO3)2(aq) + 2NO2(g) + 2H2O(l)

2. 3Pb(s) + 8HNO3(dil.) hot → 3Pb(NO3)2(aq) + 2NO(g) + 4H2O(l)

With Zinc: 1. Zn(s) + 4HNO3(conc)

hot → Zn(NO3)2(aq) + 2NO2(g) + 2H2O(l)

2. 4Zn(s) + 10HNO3(dil.) hot → 4Zn(NO3)2(aq) + NH4NO3(aq) + 3H2O(l)

5. Reaction with Some Reducing agents: Nitric acid which is a strong oxidizing agent also undergoes redox reactions with some common reducing agent like H2S, FeSO4 and SO2 etc. With H2S:

H2S(s) + 2HNO3(aq) → S(s) + 2NO2(g) + 2H2O(l)

With FeSO4: 6FeSO4(aq) + 2HNO3(aq) + 3H2SO4(aq)

→ 3Fe2(SO4)3(aq) + 2NO(g) + 4H2O(l)

With SO2: SO2(s) + 2HNO3(conc)

hot → H2SO4 + 2NO2(g)

6. As oxidizing agent): Nitric acid acts as a powerful oxidizing agent because it is an electron acceptor and has nitrogen at its highest oxidation state (+5) and can un-dergo reduction in many ways. The oxidizing property of nitric acid depends upon the following factors. (i) Concentration of the acid (ii) Nature of the reducing agent (iii) Temperature Due to these factors nitric acid gives variety of reduction products, such as NO2, NO, N2O, N2, HNO2, NH2HO3 etc.

7. Nitration Reaction (as nitrating agent): Hot conc. Nitric acid reacts with organic compounds like benzene (C6H6)

to replace hydrogen atom by the nitro (–NO2) group to form substituted

product, nitro benzene (C6H5NO2). The substitution of – NO2 group in

organic compound is called nitration.

C6H6(l) + HNO3(conc) Hot | 450°C → C6H5NO2(l) + H2O(l)

40

Q - 1) Describe Sulphur and its occurrence, uses, its physical and chem-ical properties:

SULPHUR: Sulphur is second member of VIA group of the periodic table and has symbol “S”. Its atomic number is 16 and atomic mass is 32. La-voisier recognized it as an element. OCCURRENCE OF SULPHUR: 1. Sulphur is non-metal and makes about 0.1% of the earth’s crust.

2. It is found in the Free State in Sicily, Mexico, USA, Japan New Zealand Loui-

siana and Texas.

3. In the combined state, it is largely found, as sulphide of iron, zinc, lead,

copper and mercury

4. It is also found as Sulphates of magnesium, calcium and barium.

Physical Properties of Sulphur:

1. Sulphur is a yellow solid. It is insoluble in water but soluble in carbon disulphide (CS2).

2. It is a non-metal and bad conductor of heat and electricity. 3. It melts between 113 oC to 119 oC into an amber collared liquid. 4. Its boiling point is about 444 oC. Uses of Sulphur: 1. It is used in the manufacture of Sulphuric acid, Sulphur dioxide and carbon sulphide. 2. It is used for the manufacture of calcium and magnesium hydrogen Sulphates. 3. It is used for bleaching wood – pulp. 4. It is used in volcanizing rubber and for disinfecting houses and for dirty vines. 5. It also helps to kill the fungi and insects. Chemical Properties of Sulphur: 1. Reaction with Metal: Sulphur combines with many metals directly to form their respective sulphide.

a) 2Cu(s) + S(s) heat → Cu2S(s) ← (copperas sulphide)

b) Fe(s) + S(s) heat → FeS(s) ← (Ferrous sulphide)

c) Zn(s) + S(s) heat → ZnS(s) ← (Zinc sulphide)

d) Pb(s) + S(s) heat → PbS(s) ← (Lead sulphide)

e) 2Na(s) + S(s) heat → Na2S(s) ← (Sodium sulphide)

CHEMISTRY SULPHUR AND ITS COMPOUNDS

CHAPTER

IX FOURTEEN

41

2. Reaction with Non – metal: Many non-metal react with Sulphur at different condition. With Oxygen: Sulphur burns in oxygen of the air with a bright blue flame to form Sulphur dioxide.

S(s) + O2(g) Ignition → SO2(g)

With Hydrogen: Sulphur reacts with hydrogen about 600 to 660 oC slowly to form hydrogen Sulphide.

S(s) + H2(g) 600 to 660 °C → H2S(s)

With Carbon: Sulphur combine with coke form of carbon in an electric

furnace to form a colourless liquid called carbon disulphide (CS2).

C(s) + 2S2(s) Electric furnace → CS2(l)

With Chlorine: Sulphur combine with Chlorine on heating to high tem-perature to form disulphur dichloride (S2Cl2).

Cl2(g) + 2S(s) High Temperature → S2C2

With Fluorine: Sulphur combine with fluorine on heating to form Sulphur Hexa fluoride (SF6).

S(s) + 3F2(g) Heat → SF6(l)

Q - 2) Describe and explain how Sulphur is extracted from underground deposits by Frasch process:

Frasch Process: Sulphur is extracted from its deposits deep

in the earth by Frasch Process. In this Process a hole about 30cm in diameter is drilled through the soil layers to the Sulphur bed. Structure: Three concentric iron pipes are sunk into the bore of hole. The outer most pipe is 20cm in diameter. The inner most pipe is 10cm in diameter. The middle pipe is 5cm in diameter. Procedure. Super-heated water at about 170 oC and 100 atm pressure is forced to through outer most pipe then hot compressed air at a pressure of 15atm flown down through inner most pipe to force up the molten Sulphur up through the middle pipe. The molten Sulphur is continuously pumped to solidify into wooden tanks. The Sulphur obtained is about 99.5% pure.

42

Q - 3) Describe Sulphuric Acid and its occurrence, uses, its physical and chemical properties:

SULPHURIC ACID: Sulphuric acid is one of the most important chemical com-pounds known. It is strong dibasic acid. Dilute H2SO4 dissolve some met-als liberating H2 gas. Physical Properties of Sulphuric Acid: 1. Pure Sulphuric acid is a colourless, odourless, viscous oily liquid often known

as oil of vitriol. 2. Its melting point is 10.5 0C while its boiling point is 338 0C. 3. Conc. H2SO4 (98.3%) has specific gravity about 1.84 while dilute H2SO4 is

about 1.55. 4. It is corrosive and is hygroscopic as it absorbs water vapours form surrounding

and becomes dilute. Uses of Sulphuric Acid: 1. It is used in the manufacture of fertilizers, such as di-hydrogen calcium phosphate

Ca(H2PO4)2. And ammonium Sulphate (NH4)2SO4. 2. It is used in the manufacture of Rayon, paper, plastics and detergents. 3. It is used in the manufacture of paints and pigments. 4. It is used in the making cellulose film and all kinds of main-made fabrics. 5. It is used in the steel pickling and cleaning. 6. It is used in the dehydrating, drying reactions, also oxidizing reaction. 7. It is used for refining of petroleum, motor batteries and lead accumulators. Chemical Properties of Sulphuric Acid:

Sulphuric acid behaves in three different ways. (1) As an acid (2) As an oxidizing agent (3) As drying or dehydrating agent.

1. As an acid: Sulphuric acid is a strong dibasic acid ionizes in water in two stages.

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

− ← (Hydrogen Sulphate ion)

b) 𝐻2𝑆𝑂4(𝑎𝑞)1− + 𝐻2𝑂(𝑙) 𝐻3𝑂(𝑎𝑞)

+ +𝑆𝑂4(𝑎𝑞) 2− ← (Sulphate ion)

Sulphuric acid reacts with alkalis (bases) and basic oxide like MgO to give slats and water.

1. NaOH(aq) + H2SO4((aq) → NaHSO4(aq) + H2O(l)

2. NaHSO4(aq) + NaOH(aq) → Na2SO4(aq) + H2O(l)

3. MgO(s) + H2SO4((aq) → MgSO4(aq) + H2O(l)

Hydronium ion

43

2. As an Oxidizing Agent: Sulphuric acid acts as oxidizing agent. The oxidizing properties of H2SO4 depends upon (i) con-

centration of the acid (ii) Nature of the metal or reducing agent (iii) Temperature

(i) Oxidizing of Metal:

Less electropositive metals like Zn, Fe, Al react with dilute Sulphuric to liberate H2 gas and forming their Sulphates.

a. Zn(s) + H2SO4(aq) dilute → ZnSO4(aq) + H2(g)

b. Fe(s) + H2SO4(aq) dilute → FeSO4(aq) + H2(g)

c. 2Al(s) + 3H2SO4(aq) dilute → Al2(SO4)3(aq) + 3H2(g)

Hot Conc. Sulphuric oxidizes some metals to their Sulphates liberating SO2 gas.

d. Cu(s) + 2H2SO4(aq) Hot Conc. → CuSO4(aq) + SO2(g) + 2H2O(l)

e. Pb(s) + 2H2SO4(aq) Hot Conc. → PbSO4(aq) + SO2(g) + 2H2O(l)

f. 2Al(s) + 6H2SO4(aq) Hot Conc. → Al2(SO4)3(aq) + 3SO2(g) + 6H2O(l)

Reactive metals with Conc. H2SO4 form different products. Zinc reacts

with 90% conc. H2SO4 to liberate H2S gas.

g. 2Zn(s) + 5H2SO4(g) Hot Conc. 90% → 4ZnSO4(s) + H2S(s) + H2O(l)

(ii) Oxidation of Non-metal:

Hot conc. Sulphuric acid oxidizes some non-metals like C, S and P into their oxides or oxyacid.

h. C(s) + 2H2SO4(aq) Hot Conc. → CO2(g) + 2SO2(g) + 2H2O(l)

i. S(s) + 2H2SO4(aq) Hot Conc. → 3SO2(g) + 2H2O(l)

j. 2P(s) + 3H2SO4(aq) Hot Conc. → 2H3PO4(aq) + 3SO2(g)

(iii) Oxidation with other compounds:

Conc. Sulphuric acid oxidizes hydrogen sulphide (H2S) which is oxi-

dized to Sulphur and HI is oxidized to liberate vapours of I2.

k. H2S(s) + 2H2SO4(aq) Conc. → S(s) + SO2(g) + 2H2O(l)

l. 2HI(aq) + H2SO4(aq) Conc. → I2(g) + SO2(g) + 2H2O(l)

44

3. As Drying or Dehydrating Agent: Conc. Sulphuric acid has great affinity for water, so it is able to remove hydrogen and oxygen in the form water from the compounds like sugar, ethanol, formic acid, oxalic acid etc. thus Conc. H2SO4 acts as powerful drying or dehydrating agent. This process is called dehydration.

a. C12H22O11(s)⏟ 𝑆𝑢𝑔𝑎𝑟

Hot Conc. H2SO4 → 12C(s)⏟

𝐶ℎ𝑎𝑟𝑟𝑒𝑑 𝑚𝑎𝑠𝑠

+ 11H2O(l)

b. CH3 − CH2 − OH(l)⏟ 𝐸𝑡ℎ𝑎𝑛𝑜𝑙

170 °C Conc. H2SO4 → CH2 = CH2(g)⏟

𝐸𝑡ℎ𝑒𝑛𝑒

+ H2O(l)

c. HCOOH(l)⏟ 𝐹𝑜𝑟𝑚𝑖𝑐 𝑎𝑐𝑖𝑑

Conc. H2SO4 → CO(g) + H2O(l)

d. C|C

OOH

OOH(𝑠)⏟ 𝑂𝑥𝑎𝑙𝑖𝑐 𝑎𝑐𝑖𝑑

Hot Conc. H2SO4 → CO(g) + CO2(g) + H2O(l)

e. CuSO4. 5H2O(s)⏟ Blue (hydrate)

Conc. H2SO4 → CuSO4(s)⏟

𝑤ℎ𝑖𝑡𝑒 (𝑎𝑛ℎ𝑦𝑑𝑟𝑜𝑢𝑠)

+ 5H2O(g)

Q - 4) Describe and explain how Sulphuric Acid is prepared by Contact process, what catalyst used in the process?

Contact Process: Sulphuric is prepared on large scale by contact process. This method

was developed in Germany. In this process, SO2 is produced mainly by burning Sulphur in dry air or iron pyrite in pyrite burners. Preparation of H2SO4 The given steps involves in preparation of H2SO4. 1. Preparation of SO2 2. Purification of SO2 3. Oxidation of SO2.

4. Absorption of SO3 5. Dilution of Oleum.

(i) Preparation Of SO2:

SO2 is produced mainly by burning Sulphur in dry air or iron pyrite in pyrite burners.

S(s) + O2 → SO2(g)

4FeS2(s)⏟ Iron pyrite

+ 11O2(g) → 2Fe2O3(s) + 8SO2(g)

(ii) Purification Of SO2:

SO2 contains a number of impurities such as dust particles, Arsenous oxide, vapours, Sulphur

etc. These impurities must be removed otherwise catalyst loses its efficiency (catalyst poi-

soning).

DUST CHAMBER:

SO2 is first passed through the dust chamber where steam is spread over the gas to

remove dust particles, which settle down. Fe(OH)3 also sprayed over to remove ox-

ides of Arsenic.

45

WASHING TOWER: SO2 is then passed through

a washing tower after cool-

ing. Here it is sprayed by

water to remove any other

soluble impurities. DRYING TOWER: The gas is now dried by

passing through drying tower where conc. H2SO4 (dehydrating agent) is sprayed. H2SO4

removes moisture from SO2. TYNDALL BOX (TEST BOX): Arsenic oxide is a poison for the catalyst. It is removed when the gas is passed over ferric

hydroxide. As2O3 + 2Fe(OH)3 → 2FeAsO3 + 3H2O.

In order to remove traces of As2O3, it is passed through a test box, where a strong

beam of light is thrown against the gas. If there is no scattering of light in the box, it

indicates that gas is free from As2O3.

(iii) Oxidation Of SO2 to SO3: Contact Tower: Oxidation of SO2 is carried out in contact tower where V2O5 is filled in

different pipes. SO2 here reacts with air (O2) to produce SO3. Under above conditions 98% SO2 is converted into SO3. 2SO2 + O2

→ 2SO3 + 45Kcal

Conditions necessary for maximum yield of SO3:

Oxidation of SO2 is a reversible and exothermic process in which volume of product is less than the volumes of reactants. In order to obtain maximum amount of SO3, According to Le-Chatelier’s Principle following conditions are necessary.

Concentration, Excess of O2, Temperature

A decrease in temperature favours reaction in forward direction. Optimum tempera-ture for this process is 450oC to 500oC. Pressure: Since volumes of reactants are greater than the product (3:2).Therefore, Accord-ing to Le-Chatelier’s Principle a high pressure is favourable. Optimum pressure is about 1.5 to 1.7 atmosphere. USE of catalyst: At low temperature, rate of reaction decreases. To increase rate of reaction

a catalyst Vanadium Penta-oxide (V2O5) is used.

(iv) ABSORPTION OF SO3 IN H2SO4 SO3 is not directly passed in water, because a dense fog of minute particles of H2SO4 is produced. It is therefore, dissolved in Conc. H2SO4 to form Pyro-Sulphuric acid (Oleum).

SO3 + H2SO4 → H2S2O7 ← (Oleum)

(v) DILUTION OF OLEUM Oleum is now diluted with water to form H2SO4 of required concentration.

H2S2O7 + H2O → 2H2SO4 ⇠ 𝑆𝑢𝑙𝑝ℎ𝑢𝑟𝑖𝑐 𝐴𝑐𝑖𝑑

46

Q - 5) What is allotropy? Describe different allotropic forms of Sul-phur?

Allotropy: The existence of an element in two or more different forms in the same

physical state is called Allotropy. And different forms are known as allotropic modifications or allotropes. Allotropic Forms of Sulphur: Sulphur exists in several allotropic forms, out of which the three forms are as follows: 1. Rhombic Sulphur (𝛼 − 𝑠𝑢𝑙𝑝ℎ𝑢𝑟) 2. Monoclinic Sulphur (𝛽 𝑜𝑟 𝑝𝑟𝑖𝑠𝑚𝑎𝑡𝑖𝑐 𝑠𝑢𝑙𝑝ℎ𝑢𝑟)

3. Plastic Sulphur (𝛾 − 𝑠𝑢𝑙𝑝ℎ𝑢𝑟)

(i) Rhombic Sulphur (𝛼 − 𝑠𝑢𝑙𝑝ℎ𝑢𝑟) It is the most stable crystal form of Sulphur at ordinary condition and free Sulphur exits in nature as rhombic Sulphur.

It is obtained as pale-yellow crystals, giving lemon yellow powder. Its melting point is 113 0C and its density is 2.08g/cm3 at 20 0C. It is insoluble in water but dissolve in carbon disulphide, benzene, disulphur di-chloride. Preparation:

It is prepared by slow evaporation of ordinary Sulphur solution in carbon disul-

phide or S2Cl2.

Which is filtered first to remove insoluble impurities present in Sulphur.

The filtrate on evaporation produces octahedral crystals of rhombic Sulphur.

Structure:

Rhombic Sulphur consists eight Sulphur molecules.

These molecules have single covalent bonds.

(ii) Monoclinic Sulphur (𝛽 𝑜𝑟 𝑝𝑟𝑖𝑠𝑚𝑎𝑡𝑖𝑐 𝑠𝑢𝑙𝑝ℎ𝑢𝑟) It is stable between 6 to 119 0C and hence it changes slowly to rhombic Sulphur at ordinary temperature.

It consists of dark yellow transparent needle-like crystals. On standing the crystals become opaque, brittle and lemon yellow. Its melting point is 119 0C and density is 1.96 g/cm3. It is soluble in Carbon disulphide but insoluble in water. Preparation:

It is prepared by slow cooling molten Sulphur until a crust is formed on the surface of it.

Pierce two holes through the crust and pour off the remaining molten Sulphur form inside.

On removing the rust the long needle shaped crystals of Monoclinic Sulphur is formed.

47

Structure:

Monoclinic Sulphur also consists eight Sulphur molecules. The only difference is

the shapes of the crystals.

The monoclinic Sulphur molecules unite together to give long, needle – shaped

crystals as shown in figure.

(iii) Plastic Sulphur (𝛾 − 𝑠𝑢𝑙𝑝ℎ𝑢𝑟) It is a super cooled form of Sulphur. It is the non-crystalline allotrope of Sulphur. Its boiling point is 444.6 0C and melting point is 113 0C. It is not considered as a true allotrope of Sulphur because it is unstable and re-verts to rhombic Sulphur on standing. Preparation:

It is prepared when ordinary Sulphur is heated carefully

up to its boing point about 444.6 0C then the molten Sul-

phur is poured into very cold water, a soft rubber like

mass is obtained which looks if it made of plastic material. Structure:

It has Zig –Zag arrangements of Sulphur atoms as shown

in figure. Q - 6) What is plastic Sulphur? Why is it elastic?

Plastic Sulphur: Plastic Sulphur is a soft rubber like mass which is prepared when

ordinary sulphur is heated carefully up to its boing point about 444.6 0C then the molten Sulphur is poured into very cold water. It is elastic because it is soft rubber like material. Q - 7) Explain why SO2 is purified before it is passed into the contact

tower for its oxidation to SO3?

Ans: SO2 is purified before it is passed into the contact tower for this oxidation to SO3 because at this stage it has impurities that poison the crystals. Q - 8) Describe what happens when:

(a) Sulphur is heated in the absence of air

(b) Sulphur is heated in presence of air

(a) When sulphur is heated in absence of air it changes into molten sulphur and if it

is powered in cold water it changes to plastic sulphur.

(b) When sulphur is heated in presence of air it burns with blue flame to form Sul-

phur dioxide (SO2).

48

Q - 1) What are Halogens? Why are they placed in VII A group in the periodic table? Describe the state of each member of the family and also their colours?

HALOGENS: Halogens are the elements of VII – A group in the periodic table and are reactive non-metals. They are powerful oxidizing agents because of their tendency to gain electron. They are placed in VII – A group in the periodic table because all these elements contain seven electrons in their valence shells.

Physical Properties of Halogens Element Florien (Fl) Chlorine (Cl) Bromine (Br) Iodine (I)

Atomic number 9 17 35 53

Electronic con-figuration

2, 7 2, 8, 7 2, 8, 18, 7 2, 8, 18, 18, 7

Outer shell 2S2 2P5 3S2 2P5 4S24P5 5S25P5

State at 20 0c Gas Gas Liquid Liquid

Colour Pale Yel-

low Pale Yellow

Reddish brown

Black

Atomic size (a0) 0.71 0.99 1.14 1.33

Ionic size (x–) 1.33 1.81 1.96 2.20

First ionization energy

1681 KJ/mol

1257 KJ/mol 1140 KJ/mol 1008 KJ/mol

Electron affinity –328 –349 –325 –295

Electron nega-tivity

4.0 3.0 2.8 2.5

Melting point -220 -101 -7 113

Boiling point -188 -35 59 183

X-X single bond energy

155 242 193 151

Reduction po-tential (volt)

2.87 1.36 1.07 0.56

CHEMISTRY HALOGENS CHAPTER

IX FIFTEEN

49

Q - 2) What are the sources of halogens? Describe the importance of Cl2, Br2, and I2, F2 in our daily life.

Source of Halogens: Halogens always occurs naturally as compounds with metals and

present in their negative ions. Fluoride (F–), Chloride (Cl–) Bromide (Br–) and Iodide (I–).

Chlorine, Bromine and iodine are found as halides, in sea water and in salt deposits.

Fluorine occurs in the minerals, the most widely spread compounds of fluorine are fluor-

spar (CaF2), cryolite (Na3AlF6) and flour apatite 3Ca3 (PaO4)4 CaF2.

Importance of Cl2

It is the most important substance i.e used in manufacture of Vinyl Chloride which is

extensively used in the preparation of PVC Plastic.

It is used in preparation of Hexa – Chlorocyclo Hexane, Chloro – Carbon, D.D.T (Di –

Chloro Di – Phenyl, Tri – Chloro – Ethene) etc.

It is used to disinfect water and is part of the sanitation process for sewage and

industrial waste.

During the production of paper and cloth, chlorine is used as a bleaching agent.

It is also used in cleaning products, including household bleach which is chlorine dis-

solved in water.

Chlorine is used in the preparation of chlorides, chlorinated solvents, pesticides, pol-

ymers, synthetic rubbers, and refrigerants.

Importance of Br2

The largest use for bromine is for brominated flame retardant i.e are used in TVs,

furniture, and even mobile phones.

It is most often used in analgesics, sedatives, and antihistamines.

Bromine is also widely used in both swimming pools and hot tubs as a purifier, disin-

fectant.

In agriculture, bromine acts as a very powerful and effective pesticide.

Bromine is also used in photography, and was once even used in automobiles. Importance of I2

Iodine has many commercial uses including pharmaceuticals, photographic chemi-

cals, printing inks and dyes, catalysts and animal feeds.

Iodide in small amounts is added to table salt in order to avoid thyroid disease.

It is used as mild antiseptic for cuts and scratches.

When iodine mixed with detergents, it is used as cleaning dairy equipment.

Importance of Fl2

It is used in many fluoro – chemicals, including high – temperature plastics, espe-

cially Teflon.

Hydrofluoric acid is used for etching the glass of light bulbs and in similar applica-

tions, Fluorine gas is the most reactive of all the elements and quickly attacks all metals -

steel wool bursts into flames when exposed to it.

Stannous Fluoride (SnF2) is used in toothpastes for preventing tooth decay.

50

Q - 3) (a) How is chlorine prepared in the laboratory? (b) Describe the commercial preparation of chlorine by the electrolysis of aqueous NaCl solution in Nelson’s Cell (c) what happens when chlorine re-acts with (i) Sn (ii) H2S (iii) CO (iv) P (v) FeCl2 (vi) H2O Give reac-tion and equation (d) Discuss the uses of chlorine.

Chlorine: Chlorine is the second member of halogen family i.e. of VIIA group

of the periodic table. It is most abundant and was discovered by C.W.

Scheele a Swedish chemist by the action of hydrochloric acid on manganese

dioxide (MnO2). Then name chlorine to gas was suggested by Sir Humphrey

Davy from Greek word chloros means pale green.

Physical Properties of Chlorine: 1. It is a greenish yellow gas with sharp pungent disagreeable and irritating chock-

ing smell. 2. It produces inflammation in the nose and throat and causes congestion in the

lung tissues. 3. It is fairly soluble in water and its solution in water is called chlorine water. 4. Its density is 3.21g/dm3 at STP and electronegativity is 3.0. 5. It boils at -34 0C and melts at -101 0C. Uses of Chlorine:

It is the most important substance i.e used in manufacture of Vinyl Chloride which is

extensively used in the preparation of PVC Plastic.

It is used in preparation of Hexa – Chlorocyclo Hexane, Chloro – Carbon, D.D.T (Di –

Chloro Di – Phenyl, Tri – Chloro – Ethene) etc.

It is used to disinfect water and is part of the sanitation process for sewage and

industrial waste.

During the production of paper and cloth, chlorine is used as a bleaching agent.

It is also used in cleaning products, including household bleach which is chlorine dis-

solved in water.

Chlorine is used in the preparation of chlorides, chlorinated solvents, pesticides, pol-

ymers, synthetic rubbers, and refrigerants.

Chemical Properties of Chlorine: 1. Reaction with Metal: Practically all metals combine with chlorine on heating to form their chlorides.

a) 2Na(s) + Cl2(g) heat → 2NaCl(s) ← (Sodium Chloride)

b) 2Sb(s) + 3Cl2(g) heat → 2SbCl3(s) ← (Antimony Chloride)

c) 2Fe(s) + 3Cl2(g) heat → 2FeCl3(s) ← (Ferric Chloride)

d) Zn(s) + Cl2(g) heat → ZnCl2(s) ← (Zinc Chloride)

e) Sn(s) + 2Cl2(g) heat → SnCl2(s) ← (Stannic Chloride)

51

2. Reaction with Non – metal: Many non-metal react with Chlorine at different condition. With Phosphorus (P): Phosphorus catches fire in chlorine, forming phosphorus trichloride (PCl3) a colourless liquid, in excess of chlorine forms phosphorus pentachloride (PCl5) a pale yellow solid.

2P(s) + 3Cl2(g) Heat → 2PCl3(l)

2P(s) + 5Cl2(g) Heat excess → 2PCl5(l)

With Hydrogen (H): Chlorine combine with hydrogen in the presence of sunlight or on heating quickly to form hydrogen chloride.

H2(g) + Cl2(g) Heat | Sunlight → 2HCl(g)

With Sulphur (S): Chlorine combine with hot sulphur, forming sulphur monochloride (S2Cl2), a yellowish liquid.

2S(s) + Cl2(g) Heat → S2Cl2(l) ⟵ (𝑦𝑒𝑙𝑙𝑜𝑤𝑖𝑠ℎ 𝑙𝑖𝑞𝑢𝑖𝑑)

◙ Chlorine replaces one or more atoms from other compounds such reactions are called Substitution Reaction. With Hydrogen sulphide (H2S):

H2S(g) + Cl2(g) Heat → 2HCl(g) + S(s)

With Potassium bromide (KBr):

2KBr(s) + Cl2(g) Heat → 2KCl(s) + Br2(g)

With Potassium Iodide (KI):

2KI(s) + Cl2(g) Heat → 2KCl(s) + I2(g)

With Methane (CH4):

CH4(g) + Cl2(g) HV → CH3Cl(l) + HCl

With Chloromethane (CH3Cl):

CH3Cl(l) + Cl2(g) HV → CH2Cl2(l) + HCl

3. Reaction with Ammonia: Chlorine reacts with ammonia violently to form nitrogen and hydrogen chloride, hydrogen chloride then combines with excess of NH3 to produce white fumes of NH4Cl.

2NH3(g) + 3Cl2(g) Heat → N2 + 6HCl(g)

6NH3(g) + 6HCl(g) Heat → NH4Cl ← (white fumes)

52

4. Reaction with Alkalis: When Cl2 gas is passed through caustic soda solution in cold then sodium hy-

pochlorite and chlorides are formed. ↓ (Sodium chlorate)

2NaOH(aq) + Cl2(g) Cold → NaCl(aq) + NaClO + H2O(l)

If excess of Cl2 gas is passed through hot solution of NaOH, sodium chlorate

(v) is produced, by the decomposition of sodium hypochlorite.

6NaOH(aq) + 3Cl2(g) Hot → 5NaCl(aq) + NaClO3 + 3H2O(l)

3NaClO3⏟ 𝑆𝑜𝑑𝑖𝑢𝑚 𝑐ℎ𝑙𝑜𝑟𝑎𝑡𝑒 (𝑙)

Hot | exces of Cl2 → NaClO3⏟

𝑆𝑜𝑑𝑖𝑢𝑚 𝑐ℎ𝑙𝑜𝑟𝑎𝑡𝑒 (𝑣)

+ 2NaCl(aq)

5. Reaction with Lime water: Cl2 gas react with lime water in three ways: (i) With Cold & dilute lime water in excess of Cl2 it produces calcium chlorate

(l). ↓ (Calcium hypochlorite)

2Ca(OH)2 + 2Cl2(g) Cold → Ca(OCl)2(l) + CaCl2 + 2H2O(l)

(ii) With Hot lime water in excess of Cl2 it produces calcium chlorate (v) ↓ (Calcium Chlorite V)

6Ca(OH)2 + 6Cl2(g) Hot → Ca(OCl3)2(l) + 5CaCl2 + 6H2O(l)

(iii) With Dry slaked lime water; When Cl2 is passed at over dry slaked lime at

about 40 0C it then produces bleaching powder. ↓ (Bleaching Powder)

2Ca(OH)2(𝑠)⏟ 𝐷𝑟𝑦

+ Cl2(g) 40 °C → Ca(OCl)Cl(s) + 2H2O(l)

6. As addition reaction: Cl2 gas directly combines with many compounds to form addition products.

↓ (Calcium hypochlorite)

(i)

⬚⬚CH2Ethene

⬚⬚

= CH2 + Cl2

⬚

⬚⬚

→ ⬚

Cl|

⬚Cl|

CH2 − CH2 1, 2 − dichloro ethane

(ii) CO(g) + Cl2(g) → COCl2(g) ← (Phosgene gas, a poisonous gas and

is used as chemical weapon in warfare.)

(iii) SO2(g) + Cl2(g) → SO2Cl2 ← Sulphuryl chloride

Sodium chlorate (v)

53

7. As oxidation and bleaching action: Cl2 gas in the presence of water acts as powerful oxidizing and bleaching agent due to the formation of hypochlorous acid, Cl2 reacts with water and undergo auto oxidation - reduction forming hypochlorous acid and hydrochloric acid.

Cl20 + H2O

air → H − Cl + HOCl ← hypochlorous acid

It oxidize green coloured ferrous chloride (FeCl2) to yellow coloured ferric chloride

(FeCL3) 2FeCl2(aq)(Green) + Cl2(water)

→ 2 FeCl3(aq)

Solution of chlorine in water is a strong oxidizing agent and oxidizes sulphurous

acid (H2SO3) into sulphuric acid (H2SO4). H2SO3 + Cl2 + H2O

→ H2SO4 + 2HCl

The reaction proceeds as,

(i) Cl2 + H2O → HCl + HOCl

(ii) HClO → HCl + [O]

(iii) H2SO4 + [O] → H2SO4

Preparation of Chlorine (Cl2): Chlorine is prepared by two Methods.


(i) Laboratory Method: Chlorine is usually prepared in the laboratory form hydrochloric acid, which is a convenient source of chlorine. Gently heated with oxidizing agent, such as MnO2, KMnO2, or KClO3. In this method MnO2 is taken in round bottom flask, fitted

with cork, containing a delivery tube. Concentrated hy-

drochloric acid is added in the flask over MnO2.

On gentle heating greenish yellow chlorine gas comes out which collected by the

upward displacement of air in gas through the delivery tube.

Chlorine in the laboratory can also be obtained by heating common salt (NaCl)

with conc. H2SO4 in the presence of MnO2. Reaction:

MnO2(s) + 4HCl(conc) Heat → MnCl2(aq) + 2H2O(l) + Cl2(g) ← ( Chlorine)

2KMnO4(s) + 16HCl(conc) Heat → 2KCl(aq) + 2MnCl2(aq) + H2O(l) + 5Cl2(g)

2NaCl(s) + 2H2SO4(conc) +MnO2(s) Heat Slow → Na2SO4(aq) +MnSO4(aq) + 2H2O(l) + Cl2(g)

K2Cr2O7(s) 14HCl(conc) Heat → 2KClaq + 2CrCl3(aq) + 7H2O(l) + 3Cl2(𝑔)

54

Note: Amalgam moves to another chamber

called "denuder", where it is treated with water to

produce NaOH which is in liquid state. Solid NaOH

is obtained by the evaporation of this solution.

(ii) Industrial/Commercial Preparation: (From Nelson’s Cell) Structure of Nelson’s Cell:

The cell consist of a U – shaped perforated steel vessel,

which act as Cathode, the cathode is coated inside with

a porous diaphragm of asbestos.

A carbon rod is suspended in the U – shaped vessel

which act as Anode.

The U – shaped vessel is supported in a steel tank hav-

ing an outlet for hydrogen gas.

Sodium hydroxide being formed is collected in a basin

located below the U – shaped steel tank. Working Of Cell: Purified brine solution is poured into the cell and electric current in passed through

it, chloride ions (Cl–) being negatively charged ions move towards anode where it loses electrons and liberated as chlorine gas.

Sodium ions (Na+) moved towards cathode where water splits into hydrogen gas and hydroxide ions (OH–).

The sodium ions (Na+) combine with hydroxide ions (OH–) to form sodium hydrox-ide (NaOH) which is collected in the basin.

Reaction: 2NaCl(l) 2Na(aq) + +2Cl(aq)

–

At Anode: At Cathode:

2Cl(aq) → Cl2 + 2e

−

2Na(aq) + 2e− → 2Na(s)

2Na(s) + 2H2O → 2NaOH(aq) + H2(g)

Costner Kellner’s Cell:

It is a rectangular tank of steel.

Inside of tank is lined with "ebonite".

Anode is made of titanium dipped in the saturated solution.

Flowing layer of mercury (Hg) at the bottom of tank serves

as Cathode. Working Of Cell: When electric current is passed through brine (NaCl), Na+ and Cl–

ions migrate towards their respective electrodes. Na+ ions are discharged at mercury cathode. The sodium deposited at mercury forms Sodium Amalgam (NaHg). Chlorine produced at the anode, is removed from the top of the cell. Reaction:

2NaCl(l) 2Na(aq) + +2Cl(aq)

–

At Anode: At Cathode:

2Cl(aq) → Cl2 + 2e

−

2Na(aq) + 2e− → 2Na(s)

2Na(s) + 2Hg → 2NaHg(l) ← (𝑆𝑜𝑑𝑖𝑢𝑚 𝐴𝑚𝑎𝑙𝑔𝑎𝑚)

55

Q - 4) Give the preparation of hydrogen chloride (HCl) in the labora-tory by the action of conc. H2SO4 over common salt. (b) Describe the commercial preparation of Hydrogen chloride by the direct combi-nation of H2 and Cl2 gases (c) what happens when Hydrogen chlorine reacts with (i) NaOH (ii) NaOHCO3 (iii) Pb(NO3)2 (iv) MnO2 (v) KMnO4 Give reaction and equation. (d) Discuss the uses of HCl.

Hydrochloride: The compound hydrogen chloride has the chemical formula

HCl. At room temperature, it is a colourless gas, which forms white fumes of

hydrochloric acid upon contact with atmospheric humidity.

Physical Properties of Hydrogen Chloride: 1. It is colourless gas with strong acidic odour and acidic taste. 2. It is highly soluble in water to form hydrochloric acid. 3. It is slightly heavier than air and fumes strongly in moist air. 4. Its density is 1.49 g/dm3 and It boils at -85.05 0C while melts at -114.2 0C. Uses of Hydrogen Chloride (HCl):

It is used for the pickling of iron and steel that is to remove the rust from

the metal surface.

It is used as chemical reagent in the laboratory.

It is used to remove CaCO3 deposits from sanitary wares and floors.

Chemical Properties of Hydrogen Chloride: 1. Reaction with water:

Hydrochloric acid is a strong acid and ionizes in water as. HCl(aq) + H2O(aq) 2H3O(aq)

+ + Cl(aq) –

2. Reaction with Alkalis: Hydrochloric acid reacts with alkalis like NaOH solution or KOH solution to pro-duce salt and water.

NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)

KOH(aq) + HCl(aq) → KCl(aq) + H2O(l)

3. Reaction with Ammonia: Hydrochloric acid reacts with Ammonia (NH3) to produce ammonium chloride; white fumes of NH4Cl is observed.

NH3 + 3HCl(aq) → NH4Cl(s) ← (𝒘𝒉𝒊𝒕𝒆 𝒇𝒖𝒎𝒆𝒔)

4. Reaction with less electropositive metals: Hydrochloric acid reacts with less electropositive metals like Zn, Al, Mg etc. to form hydrogen gas and their chloride. 1. Zn(s) + 2HCl(𝑎𝑞)

→ ZnCl2(aq) + H2(g) ↑ 2. Al(s) + 2HCl(𝑎𝑞)

→ AlCl2(aq) + H2(g) ↑

3. Mg(s) + 2HCl(𝑎𝑞) → MgCl2(aq) + H2(g) ↑

56

5. Reaction with metal carbonates and bicarbonates: Hydrochloric acid reacts with metal carbonates and bicarbonates to form metal chloride, carbon dioxide (CO2) and water.

Na2CO3(s) + 2HCl(aq) → 2NaCl(aq) + CO2(g) ↑ + H2O(l)

Na2HCO3(s) + HCl(aq) → NaCl(aq) + CO2(g) ↑ + H2O(l)

CaCO3(s) + 2HCl(aq) → CaCl2(aq) + CO2(g) ↑ + H2O(l)

6. Reaction with AgNO3 and Pb(NO3)2: Dil – Hydrochloric acid reacts with AgNO3 and Pb(NO3)2 to form their chlorides. And produce HNO3 (nitric acid). AgNO3(aq) + HCl(aq) (dil)

→ AgCl (s) (white precipitate) + 2HNO3(aq)

Pb(NO3)2(aq) + HCl(aq) (dil) → PbCl2(s) (white precipitate) + 2HNO3(aq)

Preparation of Hydrochloric acid (HCl): Hydrogen chloride is prepared by two Methods.


(i) Laboratory Method: The oldest laboratory method for the preparation HCl is by the action of Conc H2SO4 on common salt (NaCl). In this method NaCl is taken in round bottom flask, fitted

with cork, containing a delivery tube and thistle funnel. Con-

centrated H2SO4 is added from the thistle funnel over NaCl.

The reaction started and on gentle heating faster supply of

HCl gas is collected through upward delivery tube in gas jar.

This gas is highly soluble in water and dissolve to form hy-

drochloric acid while it is usually yellow due to impurities. Reaction:

H2SO4(conc) + NaCl(s) 150 °C

→ NaHSO4(aq) + 2HCl(g) ← ( Hydrogen Chloride)

(ii) Industrial/Commercial Preparation: (1) Hydrogen chloride is commercially prepared in the pure state by the direct com-

bination of hydrogen and chlorine gas. The reaction is strongly exothermic.

H2(g) + Cl2(g) exothermic → 2HCl(g) ; ∆H = - 44.12 KJ/mol

Sodium Hydrogen Sulphate ↓

57

(2) Hydrochloric acid from chlorides of non-metals, such as PCl3 is obtained by the action of PCl3 with water.

PCl3(l) + 3H2O(l) exothermic → H3PO3(aq) + 3HCl(g)

(3) Hydrochloric acid can also be obtained by the action of chlorine over hydrogen compounds.

H2S(g) + 2Cl2(l) → S(g) + 2HCl(g)

H2O(l) + 2Cl2(l) → O2(g) + 4HCl(aq)

Q - 5) What is bleaching powder? Give the preparation of bleaching powder by Hasenclever process (b) Describe physical and chemical properties (c) what happens when bleaching powder reacts with (i)

water (ii) excess of water (iii) H2CO3 (weak acid) (iv) Ammonia (NH3) (v)

HCl (d) Discuss the uses of Bleaching powder.

Bleaching Powder: Bleaching powder is commonly known as bleach and mixed

salt of calcium with chloride and oxychloride. And the Formula of bleaching

powder is Ca(OCl)Cl or CaOCl2.

Physical Properties of Bleaching powder: 1. It is a white amorphous powder which strongly smells like chlorine. 2. It is frequently called as Bleach. Uses of Bleaching powder:

It is used for sterilization of drinking water and disinfecting drainages and

sewers.

It is used bleaching of cotton, linen and paper pulp.

It is used for the quick predation of Cl2 gas which is a powerful oxidizing

agent.

It is also used to prepare hypochlorous acid (HClO).

Chemical Properties of Bleaching powder: 1. Reaction with water:

Bleaching powder is generally used with reasonable quantity of water. In water, it liberates Cl2 gas on reaction.

CaOCl2(s) + H2O(l) → Ca(OH)2(aq)

+ Cl2(g)

It is decomposes into CaCl2 and Ca(OCl)2 in presence of excess of water and acts as strong bleaching agent.

2Ca(OCl)Cl(s) + H2O(l) excess of H2O → CaCl2(aq)

+ Ca(OCl)2(aq)

2. Reaction with Strong acid: Bleaching powder reacts with strong acids like HCl to liberate Cl2 gas.

Ca(OCl)Cl(s) + 2HCl(aq) strong acid → CaCl2(aq)

+ Cl2(g) + H2O(l)

Phosphorous acid ↓

58

3. Reaction with CO2 in the presence of moisture (with carbonic acid): Bleaching powder reacts slowly with carbon dioxide of air in the presence of mois-ture in the air to hypochlorous acid (HClO) which oxides and destroys impurities in the surrounding atmosphere.

2Ca(OCl)Cl(s) + CO2(g) + H2O(l) slowly → CaCO3(s)

+ CaCl2(aq) + 2HOCl(aq)

4. Reaction with ammonia (NH3): Bleaching powder reacts with ammonia to evolve Nitrogen gas (N2) and produces calcium chloride and water.

3Ca(OCl)Cl(s) + 2NH3(g) slowly → N2(s)

↑ + 3CaCl2(aq) + 3H2O(l)

Preparation of Bleaching Powder (CaOCl2): Hydrogen chloride is prepared by two Methods.


(i) Laboratory Method: In the laboratory bleaching powder is prepared by shaking a small quantity of slaked

lime Ca(OH)2 with chlorine in a jar, a white powder is produced.

Reaction: Ca(OH)2(aq) + Cl2(g)

→ Ca(OCl)Cl(s) + H2O(l)

(ii) Industrial/Commercial Preparation: On large scale bleaching powder is prepared by Hasenclever process. Hasenclever Plant:

This plant consists of four iron cylinder (2 to 3 m) long.

There is an inlet in the upper most cylinder for Ca(OH)2.

The bottom cylinder has an inlet for Cl2 and outlet for bleaching powder.

Each cylinder is connected to the other by means of pipes.

Procedure:

Chlorine gas is passed in the opposite direction to come in contact with slacked lime. It is blown down form top of the cylinder by means of compressed air is allowed to fall. Result:

Chlorine gas is completely absorbed in the upper cylinders where fresh slacked lime enters. The reaction occurs to form bleaching powder (CaOCl2).

It is stored in well – ventilated rooms.

Reaction:

Ca(OH)2(aq) + Cl2(g) → Ca(OCl)Cl(s) + H2O(l)

Bleaching Powder ↓

Bleaching Powder ↓

59

Q - 6) Describe Silver Nitrate (AgNO3) test for the presence of Cl- ion in a salt of chlorine. Give the reaction.

Test for Halide irons (Silver Nitrate Test): For the detection of Cl-, Br- and I- ions silver nitrate test is performed which

gives the precipitates of AgCl, AgBr and AgI. PROCEDURE: For this purpose, the aqueous solution of the common metallic halides are treated with the solution of silver nitrate to give the precipitates of AgCl, AgBr and AgI

which helps us for the inference of Cl-, Br- and I- ions. Reactions:

AgNO3(g) + NaCl(aq) → AgCl(s) + NaNO3(g)

AgNO3(g) + NaBr(aq) → AgBr(s) + NaNO3(g)

AgNO3(g) + NaI(aq) → AgI(s) + NaNO3(g)

EXPERIMENT OBSERVATION INFERENCE

Original solution + AgNO3

1) If white ppt of AgCl soluble in dilute and conc. NH3 solution. Or

White ppt of AgCl turns purple

grey in sunlight.

Cl- is present

2) If light yellow or cream precip-itate of AgBr, which is partially

soluble in dilute NH3 solution

but insoluble in conc. NH3.

Cream ppt of AgBr turns green

yellow in sunlight.

Br- is present

3) If bright yellow ppt of AgI which is insoluble in dil. and conc.

NH3 solution. Yellow ppt of AgI

has no effect of sunlight.

I- is present

Q - 7) Identify the following: Questions Answers

a) A pale green gas that dissolves in NaOH(aq) solution give a solu-

tion used as bleach. Chlorine (Cl)

b) A gas with pungent smell, acidic taste, highly soluble in water

to form a strong acid. Chlorine (Cl)

c) White amorphous solid which reacts with water to liberate chlo-

rine. Bleaching powder

d) The deficiency of leads to enlargement of thyroid gland. Iodine (I)

e) A pale yellow gas, very unstable in water. Fluorine (F)

f) A non-sticking material, having very low coefficient of heat. Teflon or Poly tetra fluoro ethane (PTFE)

60

Q - 1) What do you mean by metal and non-metal? Also describe the physical and chemical properties of metals and non-metals.

Metals: metal is an element which ionizes by the loss of electrons to form positive ions easily.

Non – Metals: non – metal is an element which ionizes by the gain of elec-trons to form negative ions easily. Physical properties of Metal and Non – metals: (i) Metals’ Physical Properties: 1. All metals except mercury are solids with high melting points and boiling points. 2. They have characteristic lustre, known as metallic lustre and can be polished. 3. They on hitting with hammer give off notes i.e they are Sonorous. 4. They are malleable & ductile i.e they can be converted into sheets and wires. 5. They have great tensile strength and can withstand stress and strain. 6. They have relatively high densities. 7. They are good conductors of heat and electricity. (ii) Non – Metals’ Physical Properties: 1. Non – metals have low melting & boiling points about half of them are gases. 2. Non – metals do not have lustre, and cannot be polished. 3. They are not Sonorous and break on hitting. 4. They are usually brittle & break easily when subjected to stress or strain. 5. They are neither malleable nor ductile. 6. They are generally bad conductors of electricity. 7. They have relatively low densities. Chemical properties of Metal and Non – metals: (i) Metals’ Chemical Properties:

1. The oxides of metals are basic in character.

2. They turns red litmus to blue.

3. They forms positive ions during chemical reactions.

(ii) Non – Metals’ Chemical Properties: 1. The oxides of non-metals are acidic in character.

2. They turns blue litmus to red.

3. They forms negative ions during chemical reactions.

CHEMISTRY METALS AND THEIR EXTRATION

CHAPTER

IX SIXTEEN

61

Chemical reactions of Metals and Non-metals: 1. Reaction with acids:

Metals which are more electropositive than hydrogen (H) reacts with dilute acids to liberates hydrogen gas (H2) forming their salts by the loss of electrons.

1. Zn(s) + 2HCl(𝑎𝑞) → ZnCl2(aq) + H2(g) ↑ 2. Al(s) + 2HCl(𝑎𝑞)

→ AlCl2(aq) + H2(g) ↑

3. Mg(s) + 2HCl(𝑎𝑞) → MgCl2(aq) + H2(g) ↑

Non-metals do not react with dilute acids, however, with hot concentrated acids, more non-metals get oxidized into their oxides or some oxy-acids.

C(s) + 4HNO3(conc) hot → CO2(g) + 4NO2(g) + 2H2O(l)

S(s) + 6HNO3(conc) hot → H2SO4(aq) + 6NO2(g) + 2H2O(l)

S(s) + 2H2SO4(conc) hot → 3SO2(g) + 2H2O(l)

2. Reaction with Alkalis: Hydrochloric acid reacts with alkalis like NaOH solution or KOH solution to pro-duce salt and water.

NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)

KOH(aq) + HCl(aq) → KCl(aq) + H2O(l)

3. Reaction with Ammonia: Hydrochloric acid reacts with Ammonia (NH3) to produce ammonium chloride; white fumes of NH4Cl is observed.

NH3 + 3HCl(aq) → NH4Cl(s) ← (𝒘𝒉𝒊𝒕𝒆 𝒇𝒖𝒎𝒆𝒔)