WBBSE Class 10 Physical Science Notes Chapter 8.4 Inorganic Chemistry in the Laboratory and in Industry

Comprehensive WBBSE Class 10 Physical Science Notes Chapter 8.4 Inorganic Chemistry in the Laboratory and in Industry can help students make connections between concepts.

Inorganic Chemistry in the Laboratory and in Industry Class 10 WBBSE Notes

Discovery: In 1774 Priestly prepared the ammonia gas by heating a mixture of ammonium chloride and calcium hydroxide.

Occurrence: Traces of ammonia occur in the atmosphere. Ammonia is a product of the decomposition of organic matter containing nitrogen. The stable manure e.g. urea [CO(NH₂)₂], derived from the urine of animals. The urea is converted by the action of bacteria into ammonium carbonate which slowly decomposes, yielding ammonia and hence its smell is a stable and its presence in traces in air.

• Molecular formula : NH₃
• Molecular weight : 17

Preparation of ammonia:
Principle Normally ammonia gas is prepared in the laboratory by heating a mixture of ammonium chloride (1 part) and slaked lime (3 parts). Instead of using slaked lime quick lime also be used.
Chemicals required: Ammonium chloride (NH₄Cl) and quick lime (CaO) or dry slaked lime [Ca(OH)₂]

Chemical reactions:
2NH₄Cl + Ca (OH)₂ = 2NH₃↑ + CaCl₂ + 2H₂O
2NH₄Cl + CaO = 2NH₃↑ + CaCl₂ + H₂O

Collection: Ammonia is lighter than air, it may be collected by the down­ward displacement of air. Ammonia is not collected through the downward displacement of water because, it is highly soluble in water. Precautions: The ingradients, the test-tube, the delivery pipes and the gas- jar should be absolutely dry. All the connections in the apparatus should be leak-proof.

Drying of ammonia : As ammonia is a basic substance, it cannot be direct by acidic drying agents like cone. H₂SO₄ or P₂O₅.
2NH₃ + H₂SO₄ = (NH₄)₂SO₄; 6NH₃ + P₂O₅+ 3H₂O = 2 (NH₄)₃PO₄ 

It forms an additive compound with calcium chloride. So fused CaCl₂ cannot also be used to dry ammonia.
CaCl₂ + 8NH₃ = CaCl₂-8HO (Additive compound)
It is best dried with the basic drying agent, quicklime (CaO).

2. Other methods of preparation of Ammonia

(a) By hydrolysis of a metallic nitride by heating with an alkali solution.
Mg₃N₂ + 6H₂O= 3Mg (OH)₂ + 2NH₃↑
A1N + NaOH + H₂O = NaAlO₂ + NH₃↑

(b) By heating a solution of a nitrate or nitrite with zinc and strong caustic soda solution.
3NaNO₃ + 8Al + 5NaOH + 2H₂O = 8NaAlO₂ + 3NH₃↑
NaNO₂ + 3Zn + 5NaOH = 3N₂ZnO₂ + H₂O + NH₃↑

(c) By heating ammonium salts at high temperature.
(NH₄)₂SO₄ = NH₃↑ + NH₄HSO₄
2(NH₄)₃PO₄ = 6NH₃↑ + P₂O₅ + 3H₂O

(d) By the reduction of the oxides of nitrogen (except nitrous oxide) e.g. by passing a mixture of nitric oxide and hydrogen over heated spongy platinum.
2NO + 5H₂ = 2NH₂+ 2H₂O

(e) Calcium cyanamide is hydrolysed by highly heated steam to form ammonia.
CaCN₂ + 3H₂O = CaCO₃ + 2NH₃↑

(f) Harber’s process : In Haber’s synthetic process the mixture of one volume of nitrogen and three volume of hydrogen is heated at 550°C temperature and 200 atm pressure under the influence of iron as catalyst and MO as promoter to form ammonia.

Physical properties of ammonia :

• Smell Ammonia is a colourless gas with very pungent smell that affects eyes, nose and throat.
• Density Ammonia has a density (vapour density = 8 5) less than air (vapour density of air = 144). The gas is readily liquefied by pressure alone (6 atm at 0°C).
• It is highly soluble in water. 1 volume of water at 0°C dissolves 1150 volumes of ammonia and 739 at 20°C.

Liquor ammonia : A saturated aqueous solution of ammonia (sp. gr. = 0.88) contains only 35% NH₃ by weight is known as liquor ammonia. Liquor ammonia is dangerous for eyes.
NH₃ + H₂O = NH₄OH

The solution of ammonia is alkaline and turns red litmus to blue. Its basic character is due to the formation of OH^–.
NH₃+ H₂O \(\rightleftharpoons\) NH₄⁺+ OH^–
A bottle of liquor ammonia should be carefully opened after cooling in ice, as there is always a high pressure inside.

Chemical properties of NH₃

(i) Reaction with oxygen:

(a) It does not support combustion, nor does it burn in air, but in oxygen it burns with a greenish-yellow flame, forming nitrogen and steam. 4NH₃ + 3O₂ = 2N₂ + 6H₂O

(b) In presence of heated platinum gauze catalyst at 500°C – 700°C, ammonia is oxidised to nitric oxide by air or oxygen. 4NH₃ + 5O₂ = 4NO + 6H₂O

(ii) Reaction with acid : As ammonia is a base, it readily reacts with an acid to form salt.

NH₃ + HCl = NH,Cl
NH₃ + HNO₃ = NH₄NO₃
2NH₃ + H₂SO₄ = (NH₄)₂ SO₄

The reaction of NH₃ (gas) with HCl (gas) produces NH₄Cl as a white solid.
This reaction shows the formation of a solid substance from the reaction of two gaseous substances.
NH₃(gas) + HCl (gas) = NH₄Cl (solid)

(iii) Reaction with alkali metals: Ammonia reacts with alkali metals at red-heat (360°C), forming amides which are violently decomposed by water, yielding ammonia.
2Na + 2NH₃ = 2NaNH₂ + H₂↑
2K + 2NH₃ = 2KNH₂ + H₂↑
NaNH₂ + H₂O = NaOH + NH₃↑

(iv) Reaction with chlorine (non-metal): Ammonia reacts with chlorine in two ways.

(a) In excess ammonia Excess ammonia is oxidised by chlorine forming nitro­gen and is reduced itself to hydrochloric acid. Hydrochloric acid thus formed combines with ammonia producing ammonium chloride. This reaction also proves that ammonia contains nitrogen.

(b) In excess chlorine : When excess chlorine reacts ammonia forming nasent nitrogen which again combines with chlorine producing nitrogen trichloride, an oily yellow explosive compound.

(v) Reducing property: Ammonia reduces some metal oxides at high tempera­ture lo the F corresponding metals and itself gets oxidised to nitrogen gas. Ammonia gas is passed over heated black cupric oxide which is reduced to red metallic copper and ammonia is oxidised to nitrogen.

(vi) Formation of additive compounds :
Anhydrous CaCl₂,  ZnCl₂ etc. absorb ammonia to form additive compounds.
CaCl₂ + 8NH₃ = CaCl₂ .8NH₃
ZnCl₂ + 8NH₃ = ZnCl₂ .8NH₃

(vii) Reaction with carbon dioxide :
(a) Urea is formed with the reaction of ammonia and carbon dioxide at 200°C and 150 atm pressure.

(b) Ammonia is converted into ammonium sulphate with the direct use of sul­phuric acid.
2NH₃ + CO₂ + H₂O + CaSO₄ = (NH₄)₂SO₄ + CaCO₃

(vii) Reactions with salts: Salts like ferric chloride, aluminium chloride react with ammonium hydroxide forming brown precipitate of ferric hydroxide and white gelatinus precipitate of aluminium hydroxide respectively.

(ix) Formation of complex salts: In some cases the precipitated hydroxide dissolves in excess of ammonia forming a cationic complex e.g. (a) copper sulphate gives a pale blue precipitate of basic copper sulphate with ammo­nia which dissolves in excess of the precipitant forming a deep blue solution containing the complex salt, tetra-tetra-ammine cupric sulphate [Cu(NH₃)₄] SO₄.

(b) Silver nitrate solution gives a white precipitate which quickly passess into brown oxide, soluble in excess of ammonia.

(c) Ammonia produces brown colouration or precipitate in Nessler’s reagent (an alkaline solution of potassium mercuric iodide, k₂Hgl₄

Precaution to be taken to combat the effect of NH₃ leaked from industires and ammonia tanks :
Ammonia has a strong pungent smell and is highly soluble in water. It is harmful for eyes. It gives troubles in eyes when exposed in small quantities but in excess it could damage eyes permanently.

Possible measure :
(a) As NH₃ is highly soluble in cold water, it is neces­sary to spray cold water so that dispersed ammonia gas would be dissolved and minimizes its bad effect.

(b) If individual is affected then eyes should be washed readily with cold water time and again. If ammonia is leaked from ammonia tanks then the portion from which it is leaked should be kept immersed in cold water.

Identification  of Ammonia : Ammonia can be identified by the following tests:

• Ammonia is a colourless gas with a strong pungent smell.
• White fumes are obtained when ammonia gas brought in contact with a glass rod moistened, with hydrochloric acid.
• Ammonia (or its aqueous solution) forms a brown colour precipitate with Nessler’s reagent solution (an alkaline solution of K₂Hgl₂.
• Ammonia turns moist red litmus paper into blue.
• A strip of filter paper, soaked in mercurous salt solution, when exposed to ammonia gas, turns black.

Using Ammonia:
(a) Ammonia is used for the industrial preparation of nitric acid (by Ostwald process), sodium carbonate (by Solvay process).

(b) Large quantities of ammonia are used in the manufacture of fertilizers, such as-urea [CO(NH₂)₂], ammonium sulphate [(NH₄)₂SO] etc.

(c) It is also used in pharmaceutical industries and also in the preparation of smelling salt.

(d) Ammonia is used as solvent.

(e) It is used as laboratory reagent.

(f) Liquid ammonia is used as a refrigerant in ice making.

(g) Ammonia is used as cleaning agent for removing grease.

Structure of  Ammonia molecule :
Ammonia molecule has a pyramidal sharp vith nitrogen atom at the apex. The nitrogen atom of ammonia is sp³ hybridized and < H-N-H = 106°45C

(b) Hydrogen Sulphide (H₂S)

Discovery: In 1774, Shele discovered the gas sulphurated hydrogen.

Occurrence : Hydrogen sulphide is found in volcanic gases and in many hot spring waters. Its presence in the atmosphere is in very small amount. Hydrogen sulphide is also present in rotten egg as also in leather.

Molecular formula : H₂S

Molecular weight : 34

Preparation of hydrogen sulphide :

1. Laboratory preparation:

Principle: At ordinary temperature ferrous sulphide taken in a woulfe’s bottle reacts with dilute sulphuric acid (1 volume acid and 6 volumes water), hydrogen sulphide is prepared.

Chemical required . Ferrous sulphide (FeS) and dilute sulphuric acid (H₂SO₄).
Chemical reaction : FeS + H₂SO₄ = FeSO₄ + H₂ST↑
Collection : Because the gas is heavier than air, it is collected over the upward displacement of air. Hydrogen sulphide is dissolved in water, so it is not colleced over downward displacement of water. It can also be collected in downward displacement of hot water because the gas is insoluble in hot water.

Purification : The hydrogen sulphide gas can be purified by absorbing it into a suspension of magnesium oxide in water and then regenerating it by heating magnesium bisulphide [Mg(HS)₂] thus formed at 60°C.

Pure H₂S gas is obtained by heating antimony sulphide, Sb₂S₃ with pure cone hydrochloric acid.
Sb₂S₃ + 6HCl = 2SbCl₃ + 3H₂S

Drying of hydrogen sulphide :
Hydrogen sulphide may be dried by passing it through phosphorus pen oxide (P₂O₅)

(i) Anhydrous CaCl₂ is not used for drying H₂S gas.
CaCl₂ + H₂S = CaS + 2HCl

(ii) Conc. H₂SO₄ is not used for drying H₂S gas.
H₂S + H₂SO₄ = S↓ + SO₂↑ + H₂O

(iii) Quick lime also cannot be used for drying H₂
CaO + H₂S = CaS + H₂O

(iv) Hg is not used for drying H₂ But pure H₂S does not react with Hg.

N.B. Dilute sulphuric acid is used in the preparation of H₂S because it is non-volatile and dilute H₂SO₄ does not behave as an oxidising agent but cone. H₂SO₄ oxidises H₂S when sulphur is precipited.
H₂S + H₂SO₄ = S’ + SO₂↑ + 2H₂O

Nitric acid cannot be used for preparing H₂S from a sulphide for H₂S thus produced is oxidised by HNO₃ into sulphur. Dilute HCl is not used for the preparation of H₂S because HCl is a vol­atile acid.

It is mixed with hydrogen sulphide gas in such a way, separation becomes difficult.

Procautions should be taken not to inhale the gas and not to allow its prolonged contact with skin, during handling H₂S.

2. Other methods of preparation of H₂S
From metallic sulphide : Different metallic sulphides react with dilute HCl or H₂SO₄ to form H₂S.
Na₂S + 2HCl = 2NaCl + H₂S↑
CaS + H₂SO₄ = CaSO₄ + H₂S↑
Reaction of insoluble sulphide with dilute acid reacts with insoluble .sulphide As₂S₃ to form H₂S.
As₂S₃ + 12H (Zn + dil H₂SO₄) = 12 AsH₃ + 3H₂S↑

Synthetic process : Hydrogen gas and sulphur vapour are mixed with each other at 450°C in the presence of Ni-dust to produce hydrogen sulphide.

Physical properties of hydrogen sulphide :

(i) Smell : Hydrogen sulphide is a colourless gas with bad smell like rotten eggs.
(ii) Density : The density of H₂S (vapour density = 17) is higher than that of air (vapour density of air = 14 4)
(iii) Solubility in water: Hydrogen sulphide is appreciably soluble in cold water, (4-37 volumes at 0°C, 3-40 volumes at 10°C and 2-6 volumes at 20°C) but practically insoluble in hot water. The solution of H₂S in water is acidic and it turns blue litmus red.
(iv) It is poisonous when inhaled in small amounts, it causes headache and is fatal in large amounts.
(v) It liquefies at 212 2K and freezes at -190 K to a transparent solid.

Chemical properties of H₂S
(i) Combustibility : It does not support combustion, but it burns a blue flame in excess of air or oxygen, giving water and sulphur dioxide; but sulphur is deposited, if the supply of oxygen is limited.
2H₂S + 3O₂ = 2SO₂ + 2H₂O
2H₂S + O₂ = 2Sl + 2H₂O

(ii) Acidic properties: Hydrogen sulphide is a weak dibasic acid. Its aqueous solution is acidic to litmus. With sodium hydroxide (NaOH), it forms sodium sulphide (NagS) as the normal salt and sodium hydrogen sulplide (NaHS) as the acid salt.

Reducing property :

(a) Reaction with Cl₂ : When hydrogen sulphide is passed over chlorine water, chlorine is reduced to HC1, on the other hand H₂S is oxidised to sulphur.
\(\stackrel{0}{\mathrm{C}} l_2+\mathrm{H}_2 \stackrel{-2}{\mathrm{~S}}=2 \mathrm{H} \stackrel{-1}{\mathrm{C}} 1+\stackrel{0}{\mathrm{~S}} \downarrow\)

(b) Reaction with SO₂ : Hydrogen sulphide reduces sulphur dioxide to sulphur and itself oxidised to sulphur.
\(\stackrel{+4}{\mathrm{~S}} \mathrm{O}_2+2 \mathrm{H}_2 \stackrel{-2}{\mathrm{~S}}=3 \stackrel{0}{\mathrm{~S}}+2 \mathrm{H}_2 \mathrm{O}\)

(c) Reaction with HNO₃ : H₂S reduces cone. HNO_s to brown nitrogen dioxide (NO₂) and is itself oxidised to sulphur.
\(2 \stackrel{+5}{\mathrm{H}^{\mathrm{N}}} \mathrm{O}_3+\mathrm{H}_2 \stackrel{-2}{\mathrm{~S}}=2 \stackrel{+4}{\mathrm{NO}_2}+\stackrel{0}{\mathrm{~S}}+2 \mathrm{H}_2 \mathrm{O}\)

(d) Reaction with cone. H₂SO₄ H₂S reduces cone. H₂SO₄ to SO₂ and is itself oxidised to sulphur.
\(\mathrm{H}_2 \stackrel{+6}{\mathrm{~S}} \mathrm{O}_4+\mathrm{H}_2 \stackrel{-2}{\mathrm{~S}}=\stackrel{+4}{\mathrm{~S}} \mathrm{O}_2+\stackrel{0}{\mathrm{~S}}+2 \mathrm{H}_2 \mathrm{O}\)

(e) When H₂S is passed through a yellow solution of ferric chloride (FeCl₃), the salt is reduced to colourless ferrous chloride (FeCl₂) and H₂S itself on oxidation gives a precipitate of sulphur.
\(2 \stackrel{+3}{\mathrm{~F}} \mathrm{eCl}_3+\mathrm{H}_2 \stackrel{-2}{\mathrm{~S}}=2 \stackrel{+2}{\mathrm{FeCl}_2}+2 \cdot \mathrm{\textrm {HCl }}+\stackrel{0}{\mathrm{~S}} \downarrow\)

(f) H₂S redues pink solution of potassium permanganate acidified with dilute sulphuric acid to a colourless solution.

(g) H₂S also reduces potassium dichromate solution acidified with dilute sulphuric acid-the orange-red colour of the solution turns green.

(h) H₂S reduces H₂O₂ to water and H₂S is oxidised to sulphur.

(iv) Reactions with metallic salts: Hydrogen sulphide reacts with different metallic salts giving rise different coloured metallic sulphides. The basic redicals in the salts are separated from different in colours and solubilities.

(a) The following precipitates ate obtained in acid medium
CuSO₄ + H₂S = H₂SO₄ + CuSl ↓ (Black)
Pb(NO₃)₂ + H₂S = 2HNO₃ + PbSl ↓(Black)
CdCl₂ + H₂S = 2HCl + CdS ↓ (Yellow)
SnCl₂ + H₂S = 2HCl + SnS ↓ (Brown)

(b) Certain sulphides are precipitated only in alkaline solution. In ammonical solution zinc salts give a white sulphide, iron salts give black sulphide.
ZnSO₄ + (NH₄)₂S = (NH₄)₂SO₄ + ZnS ↓ (White)
FeSO₂ + (NHS = (NH₄)₂SO₄ + FeS ↓ (Black)

Identification of Hydrogen sulphide:

Hydrogen sulphide can be identified  by the following tests.

(i) It is a colourless gas with foul smell like rotten egg.
(ii) A bright silver coin, when held in the gas, turns black.
2Ag+ H₂S = Ag₂S4- (black) + H₂
(iii) Lead acetate paper turns black when it is held in H₂S gas.
Pb(CH₃COO)₂ + H₂S = PbS↓ (black) + 2CH₃COOH
(iv) The gas is passed through a solution of sodium hydroxide solution and then sodium nitroprusside solution is added to it. The solution turns purple colour.

Absorbent: An acidic gas H₂S is absorbed by the caustic alkalis, NaOH and KOH; lead nitrate solution also absorbs the gas.
Pb (NO₃)₂ + H₂S = 2HNO₃ + PbSl (black)

Uses of H₂S :

• H₂S is used as a reagent in the separation of metal ions in group analysis.
• It is sometimes used as a reducing agent.
• It is used for preparing metallic sulphides. Some sulphides are used as pigments.

Structure of hydrogen  Sulphate : It has bent structure like water and H-S-H bond angle is 92-2°. In H₂S molecule S atom is sp³ hybridized.

C. Nitrogen (N₂)

Nitrogen was discovered by Danil Rutherford in 1772 Laviosier showed that nitrogen is an elementary gas present in air and it is not a supporter of combustion and respiration.

Atomic Number : 7; Symbol: N ; Molecular formula : H₂; Atomic weight : 14008

Electronic configuration : \(\text { Is } 2 s^2 2 p_x^1 2 p_y^1 2 p_z^{11}\)
Postion in priodic Table : Period 2, Group VA ;
Oxidation number : -3 to + 5.

Comparation of Nitrogen :

A concentrated aqueous solution of ammonium nitrite is heated to produce nitrogen gas. But this reaction is very fast so there is a chance of explosion. For this reason, the concentrated aqueous solution of sodium nitrite and ammonium are mixed molar ratio then if this mixture is heated nitrogen gas is produced.

Chemical  Requird : (i) Sodium nitrite (NaNO₂)
(ii) Ammonium Chloride (NH₄Cl)

Chemical Reaction  :
(i) NaNO₂ + NH₄Cl = NH₄NH₂ + NaCl

Collection : Though nitrogen gas is slightly soluble in water, it is collected by the download displacement of water.

Drying : The gas is dried by passing through a U-tube containing cone. H₂SO₄.

2. Other methods of preparation of nitrogen :
(a) From ammonium dichromate : Ammonium dichromate on gentle heating decomposes violently, evolving nitrogen.

(b) From ammonia : N₂ is also formed by slowly passing chlorine into concen­trated ammonia, when ammonium chloride (NH₄Cl) and nitrogen (N₂) are formed.
8NH₃ + 3Cl₂ = 6NH₄Cl + N₂↑

(c) From urea Nitrogen is obtained by reaction with alkaline hypobromite solution.
CO (NH₂)₂ + 3NaOBr = CO₂↑+ 3NaBr + 2H₂O + N₂↑
From azide compounds: Spectroscopically pure nitrogen is obtained by heating barium azide at 300°C. The metal remains.

(d) From nitric acid : Moderately dilute nitric acid reacts with copper evolving nitric oxide, which when passed over the heated metal yields nitrogen.
Ba(N₃)₂ = Ba + 3N₂↑

(e) From nitric acid : Moderately dilute nitric acid reacts with copper evolving nitric oxide, which when passed over the heated metal yields nitrogen.
3Cu + 8HNO₃ = 3Cu (NO₃)₂ + 2NO + 4H₂O
2NO + 2Cu = 2CuO + N₂

(f) From air : By passing over red-hot copper filings which fix the oxygen as oxide of copper : 2Cu + O₂ = 2CuO, and nitrogen passes out. It contains about 1 percent argon.

Physical properties of Nitrogen :

• Nitrogen is a colourless gas without any smell or taste.
• Vapour density of nitrogen (14) is slightly less than that of air (14 4).
• It is slightly soluble in water (23 5 ml of N₂ is dissolved in 1 lit water at STP).

Chemical properties : It is an inert element at ordinary temperature because of very large of dissociation of the molecules, but it enters into combination with many elements at higher temperature.

(i) Reactions with non-metals:

(a) Reaction with hydrogen : At 550°C and 200 atmopsheric pressure in presence of iron catalyst reacts with hydrogen to produce ammonia. This is an industrial process of NH₃ (Haber’s process).
N₂ + 3H₂ \(\rightleftharpoons\) 2NH₃ + 22.4 Kcal

(b) Reaction with oxygen : Nitrogen combines with oxygen forming nitric oxide under the influence of electric arc at a temperature of 3000°C.
N₂ + O₂ \(\rightleftharpoons\) 2NO – 43.3 Kcal

(c) Covalent nitrides are obtained by reacting nitrogen with Boron and Silican at high temperature.
2B + N₂ = 2BN ; 6Si +4N₂= 2Si₃N₄

Nitrogen is absorbed by heated metals like Ca, Mg and Al to form nitrides which by Hydrolysis gives ammonia.
3Ca + N₂ = Ca₃N₂ ; Ca₃N₂ + 6H₂O= 3Ca (OH)₂ + 2NH₃↑
3Mg + N₂ = Mg₃N₂ ; Mg₃N₂ + 6H₂O = 3Mg(OH)₂ + 2NH₃↑
2Al + N₂= 2AlN ; AlN + 3H₂O = Al (OH)₃ + NH₃↑

(iii) Reaction with compounds : calcium carbide is heated in a current of nitrogen at temperature of 1100° C, calcium cyanamide and carbon are formed which is commercially known as nitrolim.

Significance of the presence of nitrogen in air : About 78% by volume of nitrogen is present in the air but neither plant nor animal tissues can directly absorb nitrogen from air except a few leguminous plants, such as pea, bean clover etc. There are two ways of fixation of nitrogen as

(i) By electric discharge
(ii) Bio-chemical reaction through bacteria.

(i) Electric discharge due to thundering : During electric discharge in the atmosphere nitrogen and oxygen present in air combine to form nitric oxide. This nitric oxide is then oxidised by atmospheric oxygen to produce nitrogen dioxide. Later this oxide upon mixing with water vapour or rain water forms nitric acid which falls upon our earth. Nitric acid then reacts with the bases present in the soil forming nitrate salts.

(ii) Fixation of nitrogen due to bacteria : Some micro organism and blue green algae convert nitrogen present in air to ammonia and nitric salts by chemical process.

Uses of nitrogen

• Atmospheric nitrogen is fixed in large quantities as ammonia, nitric acid and nitrolim.
• Liquid nitrogen is a refirgerant.
• Nitrogen gives an inert atmosphere in certain metallurgical operations.
• In making gas thermometers and for filling electric bulbs.
• Nitrogen is used in the preparation of different explosives.

d. Hydrogen Chloride (Hydrochloric Acid, HCl), Nitric Acid (HNO_s), Sulphuric Acid (H₂SO₄)

Hydrogen Chloride œ Hyrochioric Acid (HCl)

• Molecular weight; 365; Formula : HCl (Murlatic Acid)
• Prepared from sea-salt first; Prestly (1772)
• Devy established that it is a compound of hydrogen and chlorine.

Laboratory method of preparation of hydrogen chloride : Hydrogen chloride is obtained by heating a mixture of common salt or sodium chloride and concentrated sulphuric acid.

Reaction occurs in two steps :
(a) When the mixture is heated at 150°C – 200°C then sodium bisulphate and hydrogen chloride gas are obtained.

(b) When it is heated at 500°C then sodium sulphate and hydrogen chloride gas are produced.

• In the laboratory the reaction is performed at lower temperature because :
• At high temperature the flask used may be cracked.
• Sodium sulphate at higher temperature forms a hand crust and sticks to the glass. Its removal is very difficult.
• Collection : As dry hydrogen chloride gas is heavier than air, it is collected by downward displacement of air.
• Hydrogen chloride is highly soluble in water, so it is not collected over displacement of water.

Drying agent To remove water vapour, hydrogen chloride is passed over concentrated sulphuric acid.
P₂O₅ is not used for drying HCl gas because

(ii) CaO, NaOH, KOH are not used for drying HCl gas because :
CaO + 2HCl = CaCl₂ + H₂O
NaOH + HCl = NaCl + H₂O
KOH + HCl = KCl + H₂O

Nitric acid la not used to prepare HCl gas :

(a) Nitric acid is not used because during the preparation of HCl gas, nitric acid oxidises the produced HCl gas into Cl₂ gas.
HNO₃ + 3HCl = HOCl + 2Cl + 2H₂O

(v) Reaction with silver nitrate : An aqueous solution of hydrochloric acid (or any soluble metallic chloride) gives a curdy white precipitate of silver chloride with silver nitrate solution—the precipitate is soluble in ammonia but insoluble in nitric acid.
HCl + AgNO₃ = AgCl + HNO₃
AgCl + 2NH₃ = [Ag(NH₃)₂]Cl

(vi) Reaction with Nitric acid : The mixture of three volume of cone. HCl acid and one volume of cone. HNO₃ acid is called aqua-regia. Noble metals like gold, platinum etc. are dissolved in aquaregia.
3HCl + HNO₃ = NOCl + 2 [Cl] + 2H₂O
Au + 3[Cl] = AuCl₃ ; AuCl₃ + HCl = HAuCl₄ (soluble)
Pt + 4[Cl] = PtCl₄ ; PtCl₄ + 2HCl = H₂PtCl₆ (soluble)

Identification of hydrogen chloride and hydrochloric acid :

• HCl gas is identified by its strong choking smell.
• It fumes in moist air. Dense white fumes are produced when a glass-rod, moistened with strong ammonia solution, is held in the gas.
• HCl gas turns a moist blue litmus paper red. Hydrochloric acid also turns blue litmus paper red.
• It forms a white curdy precipitate of AgCl with colourless. AgNO₃ solution. The curdy precipitate is soluble in ammonium hydroxide solution.

Uses of hydrochloric acid :

• It is used in dyeing and calico-printing.
• It is used in the manufacture of glucose, glue and many useful metal-chlorides. HCl is used as reagents in chemical laboratories.
• It is used in preparing aqua-regia.
• It is used for washing (pickling) iron sheets before galvanization and tinning. It is used for the preparation of chlorine and chlorides.

Nitric Acid (HNO₃)

Molecular weight: 63 ; Formula : HNO₃ (aqua forties)
Prepared by distilling KNO₃ (nitre) with concentrated sulphuric acid : Glauber (1650)

Preparation :
Laboratory method of preparation : Nitric acid is prepared in the lab­oratory by heating a mixture of sodium nitrate (or potassium nitrate) and concentrated sulphuric acid (in 3 : 2 mole proportion)

Reaction occurs in two steps :
When the reaction is kept at 200°C -300°C temperature, then sodium bisul­phate or potassium bisulphate and nitric acid are produced.
NaNO₃ + H₂SO₄ = NaHSO₄ + HNO₃ KNO₃ + H₂SO₄ = KHSO₄ + HNO₃

At 800°C, sodium sulphate and nitric acid are formed in the reaction between sodium nitrate or potassium nitrate and concentrated sulphuric acid.
2NaNO₃ + H₂SO₄ = Na₂SO₄ + 2HNO₃
2KNO₃ + H₂SO₄ = K₂SO₄ + 2HNO₃

In the laboratory the reaction is performed at lower temperature because :
(a) At high temperature of about 800°C, nitric acid is decomposed to nitrogen dioxide (NO₂) and oxygen (O₂).
4HNO₃ = 4NO₂ + O₂ + 2H₂O

(b) Both HNO₃ and HCl are volatile in nature. So, they will be collected jointly in a receiver where HNO₃ may oxidise HCl into Cl₂. This difficulty is removed by using cone. H₂SO₄ (b.p. 338°C) non-volatile acid.

Pure HCl preparation: Pure hydrogen chloride is prepared by the action of water upon silicon tetrachloride:
SiCl₄ + 4H₂O = Si(OH)₄ + 4HCl↑

Hydrochloric acid is not prepared by dissolving hydrogen chloride gas in water directly because : HCl gas is highly soluble in water and the rate of dissolving of HCl gas in water is much higher than the rate of formation of HCl gas as a result there is a vacuum in the flask. To fill up the vacuum water in the beaker enters into the flask and creates explosion when it contacts with H₂SO₄.

Properties :

Physical properties :

• Colour ; HCl gas and hydrochloric acid are both colourless.
• Odour : HCl gas has a strong choking odour and the odour of hydrochloric acid is less choking.
• Solubility : It is highly soluble in water. At 0°C, 450 c.c. of HCl are dis­solved in 1c. of water.
• Boiling Point : HCl gas is easily converted to colourless liquid by applying pressure at low temperature. The liquid HCl is of boiling point -84’5°C. Liquid hydrochloride is transformed crystals at -111-4°C.
• Density : The density of cone, hydrochloric acid is 119 g/ml. The vapour density of HCI gas is 18.25.

Chemical properties :

(i) Reaction with alkali Hydrochloric acid reacts with alkalis forming salt and water.
NaOH + HCl = NaCl + H₂O
Ca(OH)₂ + 2HCl = CaCl₂ + 2H₂O

(ii) Reactions with metals : Metals lying above hydrogen in the electrochemical series react with dilute HCI forming hydrogen gas and chlorides of metals.
Mg + 2HCl = MgCl₂ + H₂↑
Fe + 2HCl = MgCl₂ + H₂↑

Noble metals, such as — gold, platinum, etc. are not reacts with the acid. Copper slowly dissolves in hot and concentrated acid and silver is slowly reacts in the presence of air only.
2Cu + 4HCl + O₂ = 2CuCl₂ + 2H₂O
4Ag + 4HCl + O₂ = 4AgCl + 2H₂O

(iii) The aqueous acid dissolves metallic oxides, hydroxides and carbonates. Hy­drogen chloride reacts with ammonia in presence of trace of moisture, forming dense fumes of ammonium chloride.
CuO + 2HCl = CuCl₂ + H₂O
NaOH + HCl = NaCl + H₂O
NH₃(g) + HCl(g) = NH₄Cl(s)
CaCO₃ + 2HCl = CaCl₂ + H₂O + CO₂

(iv) Reaction with MnO₂ : HCl is readily oxidised to chlorine by manganese dioxide or potassium permanganate.

• At high temperature glass retort may be cracked.
• At high temperature sodium sulphate (Na₂O₄) or potassium sulphate (KgSO) produced forms a hard emst and sticks to tine glass. It cannot be removed easily.

Concentrated HCl is not used in the preparation of HNO_a because:

• It is because of the fact that HCl is more volatile than HNO₃.
• During heating with HCl will be collected in the receiver, as a result HNO₃ will not be produced.
  3HCl + HNO₃ = NOCl 4- 2[Cl] + 2H₂O

Fumming nitric acid : When NO₂ is dissolved in concentrated HNO₃, it is then called fumming nitric acid. NO₂ is evolved as brown fumes from the acid. Hence, it is called fumming nitric acid. It is a strong oxidising agent. Fumming nitric acid is prepared when cone, nitric acid is distilled with As₂O₃ (arsenious oxide) or starch.

Large scale production : (Ostwald process, 1914) :
The following steps are followed :

Properties :

Physical properties:

• Colour : Pure nitric acid is a colourless liquid.
• Odour : Nitric acid has a choking smell.
• Solubility : It is highly soluble in water.
• Density : Cone. HNO₃ has a density of 142 g/ml
• Boiling point : Pure nitric acid boils at 86°C and freezes to a white solid at -42°C. The specific gravity of pure acid is T52.
• It fumes in air if it is kept opened.
•  Concentrated nitric acid is corrosive to skin.

Chemical properties:

(i) Acidic character : Aqueous solution of nitric acid is ionised to great extent, hence it is strong acid. It is a monobasic acid.
\(\mathrm{HNO}_3 \rightleftharpoons \mathrm{H}^{+}+\mathrm{NO}_3^{-}\)
HNO₃ turns blue litmus to red.

(ii) Reactions with alkali : It reacts with alkalis forming salt and water.
NaOH + HNO₃ = NaNO_s + H₂O
Ca(OH)₂ + 2HNO₃ = Ca(NO₃)₂ + 2H₂O

(iii) Reaction with metals : Nitric acid reacts all metals with the exception of gold and platinum forming different products. The actual product formed depends upon the following factors :

• nature of the metal.
• concentration of the acid.
• temperature

(a) Reaction with magnesium : Strong nitric acid on reaction with magnesium forms magnesium nitrate and nitric oxide.
3Mg + 8HNO₃ = 3Mg(NO₃)₂ + 2NO + 4H₂O
Very dilute nitric acid on treatment with magnesium and manganese liberates hydrogen.
Mg + 2HNO₃ = Mg (NO₃)₂ + H₂

(b) Reaction with iron : Hot and cone. HNO₃ makes metallic iron passive. Passive iron does not exhibit its normal chemical properties.

(c) Reaction with copper : Hot concentrated nitric acid reacts with copper metal forming copper nitrate and nitrogen dioxide.
Cu + 4HNO₃ = Cu(NO₃)₂ + 2NO₂ + 2H₂O

(iv) Reaction with AgNO_s solution and BaCl₂ solution : Nitric acid does not react with AgNO₃ solution and BaCl₂

(v) Decomposition : When nitric acid is strongly heated, it forms nitrogen dioxide and oxygen
4HNO₃ = 2H₂O + 4NO₂ + O₂

(vi) Oxidising nature: Since nitric acid has a strong tendency to give nascent oxygen it therefore, acts as a poweful oxidising agent both in the concentrated and dilute solutions. Concentrated nitric acid is generally related to nitrogen dioxide while the dilute acid is reduced to nitric oxide.
2HNO₃ (cone.) = 2NO₂ + H₂O + [O]
2HNO₃ (dilute) = 2NO + H₂O + 3 [O]

(a) Concentrated nitric acid oxidises copper turnings to copper nitrate and itself is reduced to brown coloured nitrogen dioxide gas.

(b) Concentrated nitric acid oxidises charcoal i.e. carbon to carbondioxide and itself reduced to nitrogen dioxide.

(vii) Aquaregia : A mixture to concentrated nitric acid (1 volume) and hydrochloric acid (3 volume) is called aquaregia; it dissolves gold and platinum.


(viii) A mixture of concentrated nitric acid and sulphuric acid is used in the nitration of aromatic compounds.

Identification of Nitric acid :

• Nitric acid is a colourless liquid with a choking smell.
• On heating with copper turnings, nitric acid produces broun fumes to NO₂ and copper turnings dissolve to form a blue solution of copper nitrate.
• Ring test (or a nitrate): Equal volumes of freshly prepared ferrous sulphate solution and dilute HNO₃ (dilute solution of a nitrate) are mixed together in a test tube and cooled. Cone. H₂SO₄ is now added carefully into the inner side of the test tube so as to form a heavy bottom layer; a brown ring is formed at the junction of the two liquids.

6FeSO₄ + 2HNO₃ + 3H₂SO₄ = 3Fe₂(SO₄)₃ + 4H₂O +NO
FeSO₄ + NO = FeSO₄.NO
(nitroso ferrous sulphate, brown in colour)

Uses of nitric acid :

• In the manufacture of explosives like T.N.T. (trinitrotoluene), nitroglycerine, picric acid etc.
• As a laboratory reagent.
• It is used in the manufacture of fertilizers, such as — calcium ammonium nitrate.
• In the manufacture of sulphuric acid by chamber process.
• In the manufacture of dyes, artifical silk and perfumes.
• It is used in the purification of gold and silver.
• It is used in making celluloid, rayon and other nitro cellulose products.

Pollution of air, water from Goldsmith’s workshop : Pure gold is of 24 carats. But in ornaments it is 22 carats or less. Gold ornaments are prepared by mixing requisite amount of copper with gold. When these gold ornaments are redesigned to any other form it requires breaking. Goldsmiths use nitric acid to purify gold ornaments.’ Fumes of nitric acid and brown NO₂ are formed in the process. These gases pollute air in the locality.
Cu + 4HNO₃ = Cu(NO₃)₂ + 2NO₂ + 2H₂O

Air pollution : Excess amount of NO₂ is coming from Goldsmith’s workshop is inhaled. There is an oozing of blood from lungs. Photochemical oxidant produced from NO and NO₂ have harmful effect on plant and animals.

Water pollution : Copper nitrate produced in Goldsmith’s workshop is poisonous. This compound is mixed in ponds through drains. As a result of this, fishes and plants like algae in water of this ponds may be destroyed to certain extent.

Remedy from the pollution : The works in a Goldsmith factory should Be carried out in a fume cupboard with effective device for the gaseous and particulate pollutants to escape high in the atmosphere.

Acid rain: The acidic gaseous oxides, such as – SO₂, NO₂ and CO₂ in the atmosphere in the formation of acid rain :
S + O₂ = SO₂ ; 2SO₂ + O₂ = 2SO₃; SO₃ + H₂O= H₂SO₄
NO + O₃ = NO₂ + O₂ ; NO₂ + O₃ = NO₃ + O₂
NO₃ + NO₂ = N₂O₅ ; N₂O₅ + H₂O= 2HNO₃

Acid rain affect

• Plantation and agriculture by washing out the soil nutrients.
• It also damages the building materials of houses, historic monuments and sculptures and may even destroy aquatic lives like fish.

Sulphuric Acid (H₂SO₄)

• Moleculars weight : 98; Formula : H₂O₄ (oil of vitriol)
• It is called : king of chemicals

Preparation:
Contact process :
Chemicals required : Sulphur or iron pyrities or spent sulphide, excess air, platinised asbestos or V₂O₅ (vanadium pentoxide) and water.

Principle :
(i) Sulphur or iron pyrites are burnt in aim to form SO₂.
S + O₂ = SO₂
4FeS₂ + 11O₂ = 2Fe₂O₃ + 8SO₂

(ii) The process is named as contact process because conversion of SO₂ to SO₃ carried out in presence of porous catalyst having large contact surface.

According to Le-chatlier’s principle, the conditions for the maximum yield of SO₃ are :

• Excess of O₂ (SO₂ : O₂ ratio is 2 : 3)
• High pressure (2 atm to avoid corrosion of plant)
• Low temperature (450°C)
• Presence of catalyst (finely divided platinum or V₂O_s)

(iii) SO₃ formed is absorbed into concentrated H₂SO₄ when oleum or fuming sulphuric acid gets formed.
SO₃ + H₂SO₄ = H₂S₂O₇ (oleum)

(iv) Oleum formed is diluted with water to get sulphuric acid or any desired concentration.
H₂S₂O₇ + H₂O = 2H₂SO₄

Sulphuric acid is not prepared directly by adding water to SO₃ because:
The gas forms a dense sulphuric fog or mist with water on account of the following highly exothermic reaction. H₂O + SO₃ = H₂SO₄ + 89.2 KJ

Fuming sulphuric acid or oleum : Fuming sulphuric acid or oleum is obtained when sulphur trioxide is passed over 98% sulphuric acid.
H₂SO₄ + SO₃ = H₂S₂O₇ (Pyro sulphuric acid)
Sulphuric acid can further be obtained if requisite amount of water is added to oleum.

Properties :

• Physical properties :
• Colour : It is a colourless oily liquid.
• Odour : It is an odourless liquid.
• Solubility : It is soluble in water in any proportions.
• Density : Cone. H₂SO₄ is a heavy liquid with a density 1-84 g/ml.
• Boiling point : Cone. H₂SO₄ boils at 338°C under atmospheric pressure.
• It produces severe burns on the skin.

Chemical properties :

(i) Dissociation : It is quite stable but on strong heating, it dissociates into SO₃ and H₂
\(\mathrm{H}_2 \mathrm{SO}_4 \rightleftharpoons \mathrm{H}_2 \mathrm{O}+\mathrm{SO}_3\)

(ii) Acidic character : In aqueous solutions, sulphuric acid turns blue litmus red indicating its acidic character. It inoises in two steps as :

Sulphuric acid is dibasic acid and reacts with metals, metaloxides and car­bonates etc. which are the characteristic reactions of an acid.

Reactions with alkalis :
NaOH + H₂SO₄ = NaHSO₄ (acid salt) + H₂O
2NaOH + H₂SO₄ = Na₂SO₄ (normal salt) + 2H₂O

Reaction with metals : Metals lying above hydrogen in the electrochemical series react with dilute H₂SO₄ yielding hydrogen gas.
Mg + H₂SO₄ = MgSO₄ + H ↑
Fe + H₂SO₄ = FeSO₄ + H₂↑
Hot concentrated sulphuric acid reacts with metal copper forming sulphur dioxide.
Cu + 2H₂SO₄ = CuSO₄ + SO₂ + 2H₂O

(iii) Reactions with solutions of AgNO_s and BaCl₂ : It does not react with AgNO₃ sulphuric acid gives a heavy white precipitation with BaCl₂ solution. This precipitation is in soluble in any mineral acids.
H₂SO₄ + BaCl₂ = BaSO₄ (white precipitate) ↓ + 2HCl

(iv) Oxidising action Hot concentrated sulphuric acid is an oxidising agent since it decomposes to give atomic oxygen. H₂SO₄ = H₂O + SO₂ + [O]

(a) Concentrated sulphuric acid oxidises copper metal to copper sulphate and itself reduced to sulphur dioxide.

(b) Hot and concentrated sulphuric acid oxidises carbon to carbon dioxide and itself reduced to sulphur dioxide.

(v) Dehydration action of cone. H₂SO₄ H₂SO₄ is a strong dehydrating agent and desiccating agent due to its great affinity for water.

(a) It dehydrates the white crystals of cane sugar to black mass of carbon.
C₁₂H₂₂O₁₁(sugar) + [H₂SO₄] = 12C + [11H₂O + H₂SO₄]

(b) Concentrated sulphuric acid absorbs water molecule from formic acid yielding carbon monoxide.
HCOOH + [H₂SO₄] = CO + [H₂O + H₂SO₄]

(c) Oxalic acid decomposes to produce carbon dioxide and carbon monoxide.

(d) The blue crystals of CuSO₄, 5H₂O are dehydrated to white, anhydrous CuSO₄.

Indentification of H₂SO₄

• H₂SO₄ is a heavy oilly liquid. It chars sugar and paper.
• When barium chloride solution is added to dilute solution of H₂SO₄, a heavy white precipitate is formed. The precipitate is insoluble in HCl, HNO₃ and NH₄OH solutions.
  BaCl₂ + H₂SO₄ = BaSO₄ ↓ + 2HCl

Uses of sulphuric acid :

• It is used to prepare HCl₃, HNO₃, H₃PO₄ etc
• It is used to prepare the fertilisers like ammonium sulphate, superphosphate of lime etc.
• In dyes, drugs and explosive industries for the manufacture of paints, pig­ments, dyes, drugs, picric acid and explosives like TNT.
• For refining petroleum.
• It is used for a laboratory reagent and for drying of gases.
• It is used for cleaning the surface of metals before carrying out electroplating,
• Sulphuric acid is used for the manufacture of rayon, photographic films, rubber and synthetic detergents.

Pollution due to SO₂ : SO₂, obtained from different sources, such as — ex­haust gas from motor vehicle, petroleum refining plant, extraction of metals etc. mixes in our atmosphere and creates environmental pollution.

Effect:

• SO₂ produces problems in eyes and also in lungs.
• Possibility of cancer.
• Asthama, broncrities, allergy in our body.

Stone cancer : Due to the corrosive action of SO₂, SO₃, and H₂SO₄ on marble stone i.e. on calcium carbonate are the main cause of damaging effect of historical monuments like the Tajmahal.
CaCO₃ (Marble stone) + H₂SO₄ = CaSO₄ + CO₂ + H₂O

An insoluble layer of CaSO₄ is formed over monuments. As a result once layer is formed sulphuric acid does not come in contact with marble stone. The reaction is stopped and there is no further decay of the monuments. But if this layer is removed by any means then decay continuous. It is called stone cancer.

Possible remedy :

• Banning of acid (sulphuric) factories by an ordinance.
• Excessive uses of motor vehicles should be restricted.
• Petroleum refining plant should be kept 90 km apart.

Distinction of HCl, HNO₃ and H₂SO₄ :