NITROGEN AND ITS COMPOUNDS - Chemistry Notes Form 3

• Nitrogen
  • Composition of Air
  • Preparation of Nitrogen from Air
  • Uses of Nitrogen
• Nitrous Oxide
• Nitric Oxide
• Nitrogen Dioxide
• Ammonia
  • Preparation of Ammonia
  • Chemical Properties of Ammonia
  • Physical Properties of Ammonia
  • Uses of Ammonia
• Nitric Acid 
  • Oswald Process
  • Reactions of Nitric Acid
  • Nitrates

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Nitrogen

• Nitrogen is a colourless and odourless gas, N₂, which is insoluble in water.
• Although it does not support life, it is not poisonous.
• It reacts only with difficulty with other elements, requiring either high temperatures, a catalyst, or both, in order to form compounds.
• The most important of these are ammonia and ammonium salts, certain nitrogen oxides and nitric acid and its salts.

Composition of air

• The atmosphere is the gaseous envelope which surrounds the earth.
• This gas, air, is a mixture consisting of about 78% nitrogen and 21 % oxygen.
• Water vapour is present in variable amounts (up to 5%), and so the composition of unpolluted air is normally based on the dry gas mixture.
• The figures below are percentages of the normal constituents by volume. 
  

  Gas	% by volume
  Nitrogen	78.1
  Oxygen	20.9
  Argon	0.93
  Carbon dioxide	0.035

• Nitrogen comprises about 78.1% of the earth's atmosphere and it is the source of the commercial and industrial gas.
• Traces of other gases, notably He, Ne, Kr and Xe are also found, while near cities and industrial areas, all sorts of pollutants are also found.
• Air is liquefied, and the oxygen (about 20.9%) boiled off at -183 ºC, leaving liquid nitrogen (which boils at -196 ºC) behind. This process is known as fractional distillation.
  
  
  

Preparation of Nitrogen from Air

Industrial preparation

• The chief source of free nitrogen is atmospheric air and nitrogen is usually prepared from it.
• Air free from dust, water vapour and carbon dioxide is compressed in a compression chamber for liquefaction.
• Firstly, the pressure on the air is increased to about 200 atmospheres. It is then released through a spiral into a low-pressure area, where intense cooling of the air takes place.
  
• The cold air goes upwards and further cools the spiral that brings in a fresh batch of purified air. In this way the cold air in the spiral gets progressively cooled when released.
• This procedure continues and the cooling becomes gradually more and more intense. Ultimately, the cooling becomes so great that the temperature drops to nearly -200oC. At this temperature the air condenses to form liquid air (Nitrogen becomes liquid at -196^oC).
• Liquid air is then led into a chamber, and allowed to warm up, by absorbing heat from the atmosphere.
• The boiling point of nitrogen is -196^oC; when this temperature is reached, nitrogen starts boiling and the vapours (gas) is collected and packed.
• The liquid left behind is mainly oxygen, which has a higher boiling point of 183^oC.
  

Prepararion from Ammonia and Ammonium Compounds

1. By treating excess ammonia with chlorine, ammonium chloride and nitrogen are formed.
   8NH_3(g) + 3Cl_{2 (g)} → 6NH₄Cl (s) + N_{2 (g)}
   NH₃ - Ammonia,  Cl₂ - chlorine, NH₄Cl - Ammonium chloride, N₂ - Nitrogen
   The products obtained are bubbled through water. The vapours of ammonium chloride dissolve in the water while nitrogen is collected separately.
2. By passing ammonia over heated metallic oxides like copper oxide and lead oxide.
   
   3CuO_(s) + 2NH_3(g) → 3 Cu_(s) + 3H₂O_(vap) + N_2(g)
   3PbO_(s) + 2NH_3(g) → 3Pb_(s) + 3H₂O_(vap) + N_2(g)
   NH₃ - Ammonia,  CuO - Copper (II) Oxide, Cu - Copper, H₂0- Water; N₂ - Nitrogen, PbO - Lead Oxide, Pb- Lead
3. By burning ammonia in oxygen
   Ammonia burns in oxygen to yield water vapour and nitrogen
   
   4NH_3(g) + 3O_2(g) → 6 H₂O_(l) + 2N_2(g)
   NH₃ - Ammonia, O₂ - Oxygen H₂0- Water, N₂ - Nitrogen
4. By heating a mixture of liquor ammonia with bleaching powder:
   When bleaching powder is treated with liquor ammonia, calcium chloride, water vapour and nitrogen are formed.
   3Ca(OCl)Cl_(s) + 2NH_3(l) → 3CaCl_{2 (aq)} + 3H₂O_(l) + N_2(g)
   3Ca(OCl)Cl - Bleaching powder; NH₃ - Ammonia; CaCl₂ - Calcium Chloride; O₂ - Oxygen; H₂0- Water; N₂ - Nitrogen
5. By the action of heat on ammonium compounds: (ammonium dichromate)
   Ammonium dichromate is an orange coloured crystalline substance. When heated it starts decomposing, with the evolution of heat. Sparks can be seen inside the test tube and therefore further heating is not necessary. The products of decomposition are, a green coloured solid of chromic oxide, water vapour and nitrogen gas.
   (NH4)₂Cr₂O_{7 (s)} → Cr₂O_3(s) + 4H₂O _(vap) + N_2(g)
   (NH4)₂Cr₂O₇ - Ammonium Dichromate Orange; Cr₂O₃ - Chromic Oxide Green
   However, collecting nitrogen by this method is difficult. As the reaction is accompanied by heat and light, it is quite violent. Also the green coloured fluffy chromic oxide gets sprayed all over and thrown out of the test tube. It is therefore difficult to control this reaction.

Laboratory Preparation of Nitrogen

1. Method A
   
   
   • Nitrogen can be prepared from the air as shown.
   • Air flows into the respirator and onto caustic soda which dissolves carbon dioxide gas.
   • It is then passed through a heated combustion tube containing heated copper turnings which remove oxygen.
   • Nitrogen is then collected over water.
   • Traces of noble gases present in air still remain in the final product.
2. Method B
   • Nitrogen can also be obtained by heating a mixture of sodium nitrite and ammonium chloride as shown.
   • The gas collected by this method is purer than one in method A, even though it contains water vapour which could have been removed if the gas is passed through concentrated sulphuric acid before collection.
     
   • In the laboratory, nitrogen is prepared by heating a mixture of ammonium chloride and sodium nitrite and a small quantity of water. If ammonium nitrite is heated by itself it decomposes to produce nitrogen gas. However, this reaction is very fast and may prove to be explosive.
     NH₄NO_2(s) → 2H₂O_(vap) + N_2(g)
   • For safety, a mixture of ammonium chloride and sodium nitrite approximately in the ratio of 4:5 by mass, is heated mildly with a small quantity of water.
   • The presence of water prevents ammonium chloride form subliming when heated.
   • Initially, the two substances undergo double decomposition to form sodium chloride and ammonium nitrite.
     NH₄Cl_(aq) + NaNO_2(aq) → NaCl_(aq) + NH₄NO_2(aq)
   • The ammonium nitrite so formed then decomposes to form nitrogen gas and water vapor.
     NH₄NO_2(aq) → 2H₂O_(vap) + N_2(g)
   • Nitrogen gas is collected by downward displacement of water.

Uses of Nitrogen

1. Nitrogen is used in high temperature thermometers where mercury cannot be used. This is because mercury boils at 356.7oC and hence cannot be used in such thermometers. A volume of nitrogen is enclosed in a vessel and introduced into the region of high temperature. Depending upon the temperature, expansion of the nitrogen volume takes place. Then applying the gas equation, the temperature is calculated.
2. Nitrogen mixed with argon is used in electric bulbs to provide an inert atmosphere. It helps in prevention of oxidation and evaporation of the filament of the bulb, giving it a longer life.
3. It is used to produce a blanketing atmosphere during processing of food stuff, to avoid oxidation of the food. It is also used when food is being canned, so that microorganisms do not grow.
4. It is used in metal working operations to control furnace atmosphere and in metallurgy to prevent oxidation of red-hot metals.
5. Nitrogen in the air helps as a diluting agent and makes combustion and respiration less rapid.
6. It is used by the chemical, petroleum, and paint industries to provide inactive atmosphere to prevent fires or explosions.
7. It is used in the industrial preparation of ammonia, which is converted into ammonium salts, nitric acid, urea, calcium cyanamide fertilizers etc.
8. Liquid nitrogen is used as a refrigerant for food, for storage of blood, cornea etc. in hospitals. Meat, fish etc., can be frozen in seconds by a blast of liquid nitrogen, which can provide temperatures below -196oC.
9. Liquid nitrogen is used in scientific research especially in the field of superconductors.
10. Nitrogen is essential for synthesis of proteins in plants. Proteins are essential for synthesis of protoplasm, without which life would not exist
11. Liquid nitrogen is used in oil fields, to extinguish oil fires.

Summary

Physical properties
Colour	Colourless
Odour	Odourles
Denisty Compared to air	Same as air

 

Chemical properties
Solubility in water	Slightly soluble
Burning	Does not support combustion
Moist pH paper	No reaction
Red rose petals	No reaction
Specific test	None

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Nitrous Oxide 

• Nitrous oxide (dinitrogen oxide), N₂O, is prepared by gentle heating of ammonium nitrate:
  
  NH₄NO_3(aq) → N₂O_(g) + 2H₂O_(vap)
• Nitrous oxide is a linear molecule.
• It has a boiling point of -88 ºC, and a melting point of -102ºC.
• It is colourless and has a faintly sweet smell. It is used as an anaesthetic, popularly called laughing gas.

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Nitric Oxide

• Nitric oxide, NO, may be prepared by the action of dilute nitric acid on copper:
  
  3Cu_(s) + 8HNO_3(aq) → 3Cu(NO₃)_2(aq) + 2NO_(g) + 4H₂O_(vap)
• It is a colourless gas, insoluble in water, which reacts with oxygen to form the brown gas nitrogen dioxide, NO₂:
  2NO_(g) + O_2(g) → 2NO_2(g)

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Nitrogen Dioxide

• Nitrogen dioxide, NO₂ is a planar molecule.
• It is a deep red-brown gas, which may be prepared by the action of concentrated nitric acid on copper:
  Cu_(s) + 4HNO_3(aq) → Cu(NO₃)_2(s) + 2H₂O_(vap) + 2NO_2(g)
• or by the decomposition of heavy-metal nitrates, such as lead nitrate:
  2Pb(NO3)_2(s) → 2PbO_(s) + 4NO_2(g) + O_2(g)
• At temperatures below 100 ºC, it forms dinitrogen tetroxide, N2O4:
  2NO_2(g)  ⇌ N₂O_4(g)
• Nitrogen dioxide will support combustion, as shown by the fact that a glowing splint of wood will ignite in this gas.

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Ammonia

• Ammonia is a colorless gas.
• It has a characteristic pungent odor.
• It is bitter to taste.
• Its vapor density is 8.5. Hence it is lighter than air (vapor density of air = 14.4).
• When cooled under pressure ammonia condenses to a colorless liquid, which boils at -33.4^oC.
• When further cooled, it freezes to a white crystalline snow-like solid, which melts at -77.7^oC.
• Ammonia is one of the most soluble gases in water. At 0oC and 760 mm of Hg pressure one volume of water can dissolve nearly 1200 volumes of ammonia.
• This high solubility of ammonia can be demonstrated by the fountain experiment.

Preparation of Ammonia

By Heating any Ammonium Salt with an Alkali

• In the laboratory, ammonia is usually prepared by heating a mixture of ammonium chloride and slaked lime in the ratio of 2 : 3 by mass.
• The arrangement of the apparatus is shown in the figure 6.2. As ammonia is lighter than air, it is collected by the downward displacement of air.
  
• Ammonium salt + Alkali → Salt + Water + Ammonia
  2NH₄Cl_(s) + Ca(OH)_2(s) → CaCl_2(s) + 2H₂O_(vap) + 2NH_3(g)
  NH₄Cl_(s) + NaOH_(s) → NaCl_(s) + 2H₂O_(vap) + 2NH_3(g)

Drying of Ammonia

• The drying agent used for ammonia is quick lime.
• Other drying agents such as concentrated sulphuric acid or phosphorus (V) oxide or fused calcium chloride cannot dry an alkaline gas like ammonia.
• Sulphuric acid and phosphorus (V) oxide are both acidic. They react with ammonia, forming their respective ammonium salt.

Industrial Preparation of Ammonia

Haber's Process

• Ammonia is manufactured by Haber's process using nitrogen and hydrogen (Fig.6.4).
• Reactants: Nitrogen gas -1 volume and hydrogen gas -3 volumes
  

Conditions

Temperature	Pressure	Catalyst	Promotor
450oC - 500oC	152000-684000 mm of Hg	Iron	Molybdenum

• The reaction in Haber’s process is exothermic and so external heating is not required once the reaction starts. Lowering the temperature to 450o - 500oC favours the reaction, but lowering the temperature below 450^o - 500^oC brings down the yield.
  
  
• Nitrogen is obtained in large scale from air. Air free from dust and carbon dioxide is cooled under high pressure and low temperature to about 200oC and then allowed to warm. As nitrogen has lower boiling point (-169oC) as compared to oxygen (-183^oC) it turns to gas leaving oxygen in liquid state.
• Nitrogen can also be obtained by heating ammonium nitrite (in small amounts)
  NH₄NO_2(aq) → N_2(g) + 2H₂O_(vap)

Chemical Properties of Ammonia

Combustibility

• Ammonia is neither combustible in air nor does it support combustion. However it burns in oxygen with a greenish-yellowish flame producing water and nitrogen.
  1. Burning of Ammonia in Oxygen
     The apparatus is set as shown below
     
     Firstly, when ammonia is passed through the longer tube and is made to ignite, it does not catch fire. Then oxygen is sent through the shorter tube. Now when ammonia is ignited, it catches fire and the following reaction takes place:
     4NH_3(g) + 3O_2(g) → 6H₂O_(vap) + 2N_2(g)  
     Although the products formed in the above reaction are insignificant, it is an extremely important reaction from viewpoint of industry. This is because in the presence of platinum, catalytic oxidation of ammonia can take place to give various important products.
     
     
     
  2. Catalytic Oxidation of Ammonia
     
     The platinum coil is heated at 800^oC in a combustion tube till it becomes white hot. Then ammonia and oxygen are passed through the tube. Under these conditions and in the presence of the catalyst, ammonia combines with free oxygen or oxygen of the air, to form nitric oxide and water vapour. 
     4NH_3(g) + 5O_2(g) → 6H₂O_(vap) + 4NO_(g)
     As colourless nitric oxide comes out into the air, it cools down and combines with the oxygen of the air to form reddish brown fumes of nitrogen dioxide.
     2NO_(g) + O_2(g)  → 2NO_{2(g) reddish brown}
     The formed nitrogen dioxide is converted into nitric acid in the presence of water and oxygen.
     2H₂O_(l)+ 4NO_2(g) + O_2(g)  → 4HNO_3(aq)
• The importance of the above reactions lies in the production of nitric acid, which is a very important industrial product.

Basic Nature

• Absolutely dry ammonia or pure liquefied ammonia is neutral.
• In the presence of water however, it forms ammonium hydroxide, which yields hydroxyl ions.
• As a result of this reaction, it exhibits basic nature.
  NH_3(g) + H₂O_(l) → NH₄OH_(aq)
  NH₄OH_(aq)⇌ NH₄⁺_(aq) + OH^-_(aq)
• It is a weak base and is perhaps the only gas that is alkaline in nature.
• Ammonia is an alkaline gas. When damp red litmus paper is introduced into the gas, it turns blue due to the presence of hydroxide ions as shown in the equation above.
  

Color Changes with Other Indicators

Indicator	initial Color	Final Color
Litmus	Red	Blue
Methyl Orange	Orange	Yellow
Phenophthalein	Colourless	Deep Pink

Test for Ammonia Gas

TEST METHOD	OBSERVATIONS	TEST CHEMISTRY
Strong pungent odour tested with damp red litmus	Litmus turns blue	Only common alkaline gas
conc. hydrochloric acid	White clouds with HCl fumes.	forms fine ammonium chloride crystals with HCl

Reaction of Aqueous Ammonia (NH₄OH) with Cations

TEST FOR	with aqueous sodium hydroxide	Test with aqueous ammonia	TEST CHEMISTRY
Magnesium(Mg2+)	White ppt. insoluble in excess	White ppt. insoluble in excess	Mg2+(aq) + 2OH-(aq) → Mg(OH)2(s) white ppt. The pp t. is not soluble in excess of NH3 or NaOH. You could distinguish Mg2+ from Ca2+ with a flame test
Calcium (Ca2+)	White ppt. insoluble in excess	No ppt. or very slight white ppt insoluble in excess	Ca2+(aq) + 2OH-(aq) → Ca(OH)2(s) White ppt. The ppt. is not soluble in excess of NH3 or NaOH.
Aluminium(Al3+)	White ppt. soluble in excess giving a colourless solution	White ppt. insoluble in excess	Aluminium ion: Al3+(aq) + 3OH-(aq) → Al(OH)3(s) in excess NaOH forms soluble (Al(OH)4-)
Zinc (Zn2+)	White ppt. soluble in excess giving a colourless solution	White ppt. soluble in excess giving a colourless solution	Zinc ion: Zn2+(aq) + 2OH-(aq) → Zn(OH)2(s) white ppt. The ppt. dissolves in both excess sodium hydroxide and ammonia to give a clear colourless solution.
Lead (Pb2+)	White ppt. soluble in excess giving a colourless solution	White ppt. insoluble in excess
Iron(II) (Fe2+)	Green ppt. insoluble in excess	Green ppt. insoluble in excess	Iron(II) ion: Fe2+(aq) + 2OH-(aq) → Fe(OH)2(s) dark green ppt. The ppt. is not soluble in excess of NH3 or NaOH.
Iron (III) (Fe3+)	Red-brown ppt insoluble in excess	Red-brown ppt insoluble in excess	Iron(III) ion: Fe3+(aq) + 3OH-(aq) → Fe(OH)3(s) brown ppt.* The ppt. is not soluble in excess of NH3 or NaOH.
Copper (Cu2+)	Light blue ppt. insoluble in excess	Light blue ppt. soluble in excess giving a deep-blue solution

Ammonia as a reducing agent

• Heated dry ammonia gas can reduce copper (II) oxide to pure copper. This reaction can be used to prepare nitrogen.
  
  NH_3(g) + CuO_(s)  → N_2(g) + H₂O_(g)
• The gas passes through a U- tube surrounded by cold water which contains some melting ice. This helps to condense the vapour produced to liquid water. Nitrogen is finally collected by downward displacement of water.

Fountain experiment

• Fill a clean dry round-bottomed flask with dry ammonia, close it by a one holed stopper, through which a long jet tube is introduced. The free end of the tube is dipped into a trough of water as shown.
  
• Add two or three drops of an acid and a small quantity of phenolphthalein to the water in the trough. This water is colorless. Pour a small quantity of spirit or ether on a layer of cotton and place it over the inverted flask.
• Due to the cooling effect produced by the process of evaporation of spirit or ether, the ammonia gas contracts a little and as a result, small quantity of the water gets sucked up. As soon as this water enters the flask, the ammonia dissolves in it, forming a partial vacuum. As a result of it, water rushes in and comes out of the tube as a jet of fountain. The color of the water turns deep pink.

Physical Properties of Ammonia

• Colourless gas
• Distinctive pungent smell
• Less dense than air
• Very soluble in water to give an alkaline solution
  
  

Dissolving Ammonia in Water

• Due to its high solubility, ammonia cannot be passed through water like many other gases. Ammonia is dissolved in water, as shown below.
  
• This arrangement is called funnel arrangement and its principle is the same as that discussed for HCl gas.
• The funnel arrangement prevents back suction of water, which can cause damage to the apparatus used.
• It provides larger surface area for dissolution of ammonia. A very strong solution of ammonia in water is called liquor ammonia. Ammonia can be obtained from it by boiling.

Action with Acids

• Ammonia reacts with the acids to form their respective ammonium salts.
• The ammonium salts appear as white fumes
• Ammonia gas + acid → ammonium salt
  NH_3(g) + HCl_(aq) → NH₄Cl_(aq)
  
  NH_3(g) + H₂SO_4(aq) → (NH₄)₂SO_4(aq)
  NH_3(g) + HNO_3(aq) → NH₄NO_3(aq)

Uses of Ammonia

The following are the chief uses of ammonia:

1. Ammonia is used in the industrial preparation of nitric acid by Ostwald's process.
2. Fertilisers, such as ammonium sulphate, ammonium nitrate, ammonium phosphate, urea etc. are manufactured with the help of ammonia.
3. It is used in the manufacture of other ammonium salts, such as ammonium chloride, ammonium carbonate, ammonium nitrite etc.
4. It finds use in the manufacture of nitrogen compounds such as sodium cynamide, plastics, rayon, nylon, dyes etc.
5. It is used in the manufacture of sodium carbonate by Solvay's process. (Ammonia and carbon dioxide are treated with aqueous sodium chloride, crystals of sodium hydrogen carbonate are formed. They are heated to yield sodium carbonate)
6. Ammonia acts as refrigerant in ice plants. Evaporation of a liquid needs heat energy. About 17g of liquid ammonia absorb 5700 calories of heat from the surrounding water. This cools the water and ultimately freezes it to ice.
7. Ammonia is used to transport hydrogen. Hydrogen is dangerous to transport, as it is highly combustible. So it is converted to ammonia, liquefied, transported and then catalytically treated to obtain hydrogen.
8. Many ammonium salts are used in medicines. Inhaling the fumes produced by rubbing ammonium carbonate in the hands can revive people who have fainted.
9. It is used as a cleansing agent. Ammonia solution emulsifies fats, grease etc. so it can be used to clean oils, fats, body grease etc. from clothes. It is also used to clean glassware, porcelain, floors etc.
10. It is used as laboratory reagent.

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Nitric Acid

• Nitric acid is produced industrially from ammonia by the Oswald process. It is a strong acid, converting bases to salts called nitrates:
  CuO_(s) + 2HNO_3(aq) → Cu(NO₃)_2(s) + H₂O_(vap) 
  NaOH_(s) + HNO_3(aq)  → NaNO_3(s) + H₂O_(vap)
• Nitric acid is also a strong oxidizing agent and may be reduced to nitric oxide or nitrogen dioxide:
  Cu_(s) + 4HNO_3(aq)  → Cu(NO₃)_2(s) + 2H₂O_(vap) + 2NO_2(g)
  3Cu_(s) + 8HNO_3(aq) → 3Cu(NO₃)_2(s) + 2NO_(g) +4H₂O_(vap)
• Pure nitric acid slowly decomposes to form water, nitrogen dioxide and oxygen. This causes the nitric acid to become yellow. The process is accelerated on heating:
  4HNO_3(aq) → 2H₂O_(l) + 4NO_2(g) + O_2(g)

Oswald Process

• Nitric acid is prepared in large scale from ammonia and air (Fig.6.14).

Reactants
Pure dry ammonia (1 volume) and air (10 volumes)
Reactions

1. 1st step - Catalytic oxidation of ammonia to form nitric oxide
   
2. 2nd step - Oxidation of nitric oxide to nitrogen dioxide.
   
3. 3rd step - Absorption of nitrogen dioxide in water to give nitric acid.
   4NO_2(g) + 2H₂O_(l) + O_2(g) → 4HNO_3(aq)

Catalyst

Platinum (for oxidation of NH₃)

Temperature
700^oC- 800^oC

Reactions of Nitric Acid

• Cuprous Oxide, Cu2O reacts with dilute Nitric Acid, HNO3, in the cold to form a solution of Cupric Nitrate, Cu (NO3)2, and Copper, Cu.
  Cu₂O_(s) + 2HNO_3(aq)  → Cu(NO₃)_2(s) + Cu_(s) + H₂O(l)
• Cuprous Oxide, Cu2O reacts with concentrated Nitric Acid, HNO3, or with dilute Nitric Acid, HNO3, on heating, when the Cuprous Oxide, Cu2O dissolves with evolution of Nitric Oxide, NO.
  3Cu₂O_(s) + 14HNO_3(aq)  → 6Cu(NO₃)_2(s) + 2NO_(g) + 7H₂O_(l)
• Dinitrogen Pent oxide, N₂O₅, is best prepared by dehydrating concentrated Nitric Acid, HNO₃, by Phosphorus Pent oxide, P₂O₅.
  2HNO_3(aq) + P₂O_5(s)  → N₂O_5(s) + 2HPO_3(aq)
• Nitric Oxide, NO is prepared by the action of Copper, Cu, or Mercury, Hg, on dilute Nitric Acid, HNO₃, and was called Nitrous Air.
  3Cu_(s) + 8HNO_3(aq)  → 3Cu(NO₃)_2(s) + 2NO_(g) + 4H₂O_(l)
• Nitrogen dioxide, NO₂, is a mixed acid anhydride and reacts with water to give a mixture of nitrous and nitric acids.
  2NO_2(g) + H₂O_(l)  → HNO_2(aq) + HNO_3(aq)
• If the solution is heated the nitrous acid decomposes to give nitric acid and nitric oxide.
  3HNO_2(aq)  → HNO_3(aq) + 2NO_(g) + H₂O_(l)
• Sulphur Dioxide, SO₂, and Nitrogen Oxides, NOx, are toxic acidic gases, which readily react with the Water, H₂O in the atmosphere to form a mixture of Sulphuric Acid, H₂SO₄, Nitric Acid, HNO₃, and Nitrous Acid, HNO₂, .
• The dilute solutions of these acids which result give rain water a far greater acidity than normal, and is known as Acid Rain.
• Nitrates are the salts of nitric acid, and are strong oxidising agents.
• The Oswald Process is the tree stage process by which Nitric Acid, HNO₃, is manufactured. Firstly, Ammonia, NH₃, is oxidised, at high temperature (900 ⁰C.) over a platinum-rhodium catalyst, to form Nitrogen Monoxide, NO.
  4NH_3(g) + 5O_{2 (g)} → 4NO_(g) + 6H₂O_(l)
• The Nitrogen Monoxide, NO, cools and reacts with oxygen, O₂, to produce Nitrogen Dioxide, NO₂.
  2 NO_(g) + O_2(g) → 2NO_2(g)
• Finally, the Nitrogen Dioxide, NO₂ reacts with Water and Oxygen, O₂, oxygen to produce Nitric Acid, .
  4NO_2(g) + 2H₂O_(l) + O_2(g) → 4HNO_{3 (l)}
  Cu₂O_(s) + 2HNO_3(aq)  → Cu(NO₃)_2(s) + Cu(s) + H₂O_(l)
• Cuprous Oxide, Cu₂O reacts with concentrated Nitric Acid, HNO₃, or with dilute Nitric Acid, HNO₃, on heating, when the Cuprous Oxide, Cu₂O dissolves with evolution of Nitric Oxide, NO.
  3Cu₂O_(s) + 14HNO₃  → 6Cu(NO₃)_2(s) + 2NO_(g) + 7H₂O_(l)

Nitrates

• Salts of metals with nitric acid are called nitrates. Most nitrates are soluble in water.
• The nitrates of alkali metals form nitrites when strongly heated:
  2NaNO_3(s) → 2NaNO_2(s) + O_2(g)
• The nitrate of other metals decompose on heating to form nitrogen dioxide and the metal oxide, or, in the case of some metals such as silver and gold, the pure metal, nitrogen dioxide, and oxygen:
  2Pb(NO₃)_2(s) → 2PbO_(s) + 4NO_2(g) + O_2(g)
  2AgNO_3(s)  → 2Ag_(s) + 2NO_2(g) + O_2(g)

Summary of Action of Heat on Nitrates

Generally compounds of very reactive metals such as sodium and potassium are more stable to heat than the metals lower down in the reactivity series of metals.

Reactivity Series for Metal	Action of heat on nitrate of metal
K Na Cu	Decompose to metal nitrite + oxygen
Ca Mg Al Fe	Decompose to metal oxide + oxygen + nitrogen dioxide
Hg Ag Au	Decompose to pure metal + Oxygen + nitrogen dioxide

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