Notes (Chapter # 13) Nitrogen and Oxygen

 UNIT# 13

NITROGEN AND OXYGEN

 

Q#	Question	Year
Q# 01	Describe four types of normal oxides	[2012]
Q# 02	Explain oxides	[2018][2017][2015][2014][2011]
Q# 03	How is ammonia manufactured by Haber Bosch process	[2011]
Q# 04	Describe in three ways Oxidation	[2017] [2015]
Q# 05	Describe three ways Reduction	[2017]
Q# 06	Write note on ozone	[2017][2010] [2005]
Q# 07	Write three differences between oxidation and reduction	[2013][2010]
Q# 08	Discuss industrial preparation of ammonia	[2018][2016] [2009]
Q# 09	Discuss industrial preparation of nitric acid	[2017][2015]
Q# 10	Explain Nitric acid by Ostwald’s method	[2013][2005]
Q# 11	What is aqua regia?	[2016][2011]

 

Q. What do you know about Nitrogen?

NITROGEN

GROUP

Nitrogen belongs to V A group in the periodic table and first member of its family containing 5 electron in its valence shell

N = 7 = K², L⁵

 

OCCURRENCE

Nitrogen occurs in the free state of N₂ gas in air up to 78% by mass of the earth’s atmosphere. In combine state, it occurs in the earth crust as nitrate of Sodium, Calcium and potassium. It is also found in combined state in organic matter such as protein, urea and vitamin B compound.

Q. Describe the preparation of Nitrogen by commercial method.

PREPARATION OF NITROGEN

COMMERCIAL METHOD

The only important commercial method of producing nitrogen gas is the fractional distillation of liquid air

 

PROCESS

1.     In this process air is liquefied to form liquid air, which is then fractionally distilled.

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

3.     The liquid air is filtered to remove CO₂ solid and then distilled.

4.     Nitrogen is the most volatile component with boiling point -196^o C distills over.

5.     Argon which boils at -185^o C is removed from the middle of the column and oxygen the least volatile component with boiling point -183^o C collects at the bottom of the column.

 

Q. Describe the preparation of Nitrogen by laboratory method.

LABORATORY METHOD

Pure nitrogen in the laboratory is prepared by heating ammonium nitrate which thermally decomposes to give nitrogen gas.

Ammonium nitrate is first obtained by reacting ammonium chloride with sodium nitrate.

NH₄Cl + NaNO₂       ⟶       NH₄NO₂ + NaCl

NH₄NO₂                          ⟶       N₂ + 2H₂O

 

Q. Describe the physical properties of Nitrogen.

PHYSICAL PROPERTIES OF NITROGEN

1.     Nitrogen is colorless, tasteless and odorless.

2.     Pure nitrogen is slightly lighter than air.

3.     It is only slightly soluble in water.

4.     Its boiling point is -196^o C.

5.     While melting point is -210^o C.

 

 

Q. Describe the chemical properties of Nitrogen.

CHEMICAL PROPERTIES OF NITROGEN

 

1)   Reaction with metal

Nitrogen combines directly with some metals like Ca, Mg, Fe, Al to form respective nitrides at high temperature and pressure.

3Mg + N₂       ⟶       Mg₃N₂

3Ca + N₂         ⟶       Ca₃N₂

 

2)   Reaction with non-metal

N₂  + O₂                            ⟶       2NO  

N₂ + 3H₂                         ⟶       2NH₃

 

Q. Describe the uses of Nitrogen.

USES OF NITROGEN

1.     Nitrogen is used in commercial preparation of ammonia and nitric acid.

2.     Compound of nitrogen are used in fertilizers industry.

3.     It is inert gas, used to provide inert atmosphere for the production of light bulb and electronic component.

 

Q. What do you know about Oxygen?

OXYGEN

Oxygen belongs to VI A group of the periodic table. It is the first member of this family. It contains 6 electrons in the valence shell.

O = 8 = K², L⁶

 

OCCURRENCE

The most abundant element on earth is Oxygen. It occur in nature both in the free and combined state.

Free State in Free State oxygen is present as diatomic gas (O₂) in the earth atmosphere. In air up to 21% by volume and 33% dissolved in water.

 

Combined State

·        It constitutes about 88.8% by mass of water (H₂O).

·        It is present in silica (SiO₂) silicates, carbonates and oxides of both metal and non-metals.

·        Human body is made up of 2/3 of oxygen.

Q. Describe the isolation of oxygen.

ISOLATION OF OXYGEN

Oxygen is isolated by the fractional distillation of liquid air just like nitrogen, as we know oxygen forms 21% by volume by air.

 

From air

The isolation of oxygen from air involves two steps, which are:

·        Liquid of air

·        Fractional distillation of liquid air

 

LIQUEFACTION OF AIR

Air in the gaseous form is first passed through caustic soda to remove CO₂ present in air. It is then compressed, under high pressure about 200 atmosphere in the compressor, then cooled and allowed to expand rapidly through a nozzle. The sudden expansion of air into a region of lower pressure causes the air to cool even further. The process of compression and expansion are repeated over and over again due to which temperature falls up to 200^oC at which air liquefies.

 

FRACTIONAL DISTILLATION OF LIQUID AIR

The liquid air is then led to a fractionating column through a filter in order to remove the traces of CO₂, solid is left behind. On distillation nitrogen with lower boiling point of 196^oC, involves first leaving behind a liquid very rich in oxygen.

Further heating turns liquid oxygen into gas which boils out at 185.7^oC and passes off from the middle of the column and liquid oxygen, the last volatile component in the air turns into oxygen gas at 183^oC. Oxygen gas is dried, compressed and stored in steel cylinders under a pressure about 100 atmospheres.

 

Q. Describe the laboratory preparation of Oxygen.

LABORATORY PREPRATION

Oxygen is prepared b heating a mixture of potassium chlorate and manganese dioxide. MnO₂ act as catalyst which cause decomposition of potassium chlorate at low temperature

2KClO₃                             ⟶       2KCl + 3O₂

 

 

 

Q. Describe the physical properties of Oxygen.

PHYSICAL PROPERTIES OF OXYGEN

1.     Oxygen is colorless, tasteless and odorless gas.

2.     It is neutral to moist litmus paper.

3.     Gaseous oxygen is about 1.1 times denser then air.

4.     It is slightly soluble in water, only 2% of it dissolves by volume at room temperature. Its solubility is of vital importance to aquatic life.

5.     Gaseous oxygen liquefies at -183^o C and solidifies at -225 ^OC

 

Q. Describe the chemical properties of Oxygen.

CHEMICAL PROPERTIES OF OXYGEN

1)   Reaction with metals

Li, Na, K of group IA Mg, Ca, and Ba of group IIA react with oxygen to give their respective oxides.

4Li + O₂                           ⟶       2Li₂O

4Na + O₂                        ⟶       2Na₂O

4K + O₂                            ⟶       2K₂O

2Ca + O₂                         ⟶       2CaO

2Mg + O₂                       ⟶       2MgO

 

2)   Reaction with non-metals

Oxygen reacts with non-metal to produce following products

S + O₂                                 ⟶       SO₂      

C + O₂                                                ⟶       CO₂

4P + 5O₂                        ⟶       P₄O₁₀

 

3)   Reaction with methane

CH₄ + 2O₂                     ⟶       CO₂ + 2H₂O

 

4)   Reaction with other compound

2H₂S + 3O₂                  ⟶       2SO₂ + 2H₂O

4FeS + 7O₂                   ⟶       2Fe₂SO₃ + 4SO₂

 

Q. Explain Oxides.

OXIDES

The binary compound of oxygen with metals and non-metal are called as “Oxides”.

CLASSIFICATION

Oxides are classified into various groups on the basis of valence no. or oxidation state of oxygen.

1.     Normal oxides

2.     Per oxides

3.     Super oxides

4.     Sub oxides

 

1)   NORMAL OXIDES

Oxides in which oxygen show normal oxidation state 2, are known as normal oxides. It is further divided into four types on the basis of their chemical characteristics. 

a.      Basic oxide

b.     Acidic oxide

c.      Amphoteric oxide

d.     Neutral oxides

 

a)   BASIC OXIDES

Basic oxides are generally the oxides of metals. They are generally ionic oxides and are white solid.

4Na + O₂                        ⟶                   2Na₂O

2Pb + O₂                         ⟶                   PbO

2Ca + O₂                         ⟶                   2CaO

Most of basic oxides dissolve in water and produce their hydro-oxide and turn red litmus blue.

Na₂O + H₂O                 ⟶                   2NaOH

CaO + H₂O                    ⟶       Ca(OH)₂

They also react with acid to form salt and water.

MgO + 2HCl                ⟶       MgCl₂ + H₂O

CaO + 2HNO₃            ⟶       Ca(NO₃)₂ + H₂O

 

b)   ACIDIC OXIDES

The normal oxides of non-metal are generally acidic. e.g.

S + O₂                                 ⟶       SO₂

C + O₂                                ⟶       CO₂

N₂ + 2O₂                         ⟶       2NO₂

 

These oxides are soluble in water to form acids which turns blue litmus red.

SO₂ + H₂O                    ⟶       H₂SO₃               

CO₂ + H₂O                     ⟶       H₂CO₃

 

They form salt and water.

CO₂ + 2NaOH            ⟶       Na₂CO₃ + H₂O

SO₃ + 2KOH                 ⟶       K₂SO₄ + H₂O

 

c)    AMPHOTERIC OXIDES

Oxygen react with less electropositive metal forms oxides that possess dual nature i.e. acidic as well as basic. These oxides known as amphoteric oxides.

4Al + 3O₂                      ⟶       2Al₂O₃

2Zn + O₂                         ⟶       2ZnO

 

Amphoteric oxides react with acids behaving just like bases to form salt and water.

Al₂O₃ + 6HCl              ⟶       2AlCl₃ + 3H₂O

ZnO + H₂SO₄               ⟶       ZnSO₄ + H₂O

 

Amphoteric oxide reacts with bases behaving just like acids, to form salt and water.

Al₂O₃ + 2NaOH        ⟶       2NaAlO3+H₂O
                                                (Sodium Aluminates)

ZnO+2NaOH              ⟶       Na₂ZnO₂ + H₂O

 

d)   NEUTRAL OXIDES

Neutral oxides are neither basic nor acidic. They are neutral to litmus. i.e. water (H₂O), Nitric acid (NO), Carbon mono oxide (CO) and nitrous oxide.

 

2)   PER OXIDES

Oxides containing higher proportion of oxygen as compared to normal oxides are per oxides. These oxides oxygen has oxidation state -1. They contain per oxide ion (O-O)²

FOR INSTANCE

·        Sodium peroxide (Na₂O₂)

·        Barium peroxide (BaO_2­)

·        Hydrogen peroxide (H₂O₂)

 

Na₂O₂ + 2HCl         ⟶       2NaCl + H₂O₂

 

3)   SUPER OXIDES

In super oxides, the valence number of electron is -1/2. They have greater amount of oxygen then normal and per oxides.

The element of group IA potassium, rubidium and cesium form super oxides. They release O₂ o heating and are powerful oxidizing agent.
e.g. KO₂, RbO₂, C₅O₂

 

4)   SUB OXIDES

They have less quantity of oxygen then the normal oxides. They are unstable. Few sub oxides are also known as Carbon sub oxide. i.e. C₂O₂

 

Q. Define the following terms.

·        Oxidation

·        Reduction

·        Oxidizing agent

·        Reducing agent

·        Redox reaction

 

OXIDATION

The reaction in which the loss of electron occurs is known as oxidation.

 

REDUCTION

The type of reaction in which gain of electron occur is known as reduction.

 

OXIDIZING AGENT

A substance which accepts or gains electron is defined as oxidizing agent and itself get reduced.

 

REDUCING AGENT

A substance that losses or donates electron is known as reducing agent.

 

REDOX REACTION

Oxidation and reduction occur simultaneously, so they are termed as redox reaction.

Redox reaction involve a transfer of electron. There is a simultaneously loss and gain of electron. The substance which electron called reducing agent and itself oxidized.

The substance accepting electron is called oxidizing agent and itself get reduced.

EXAMPLE

Rusting

Fe + 3O₂ + 2H₂O                    ⟶       2Fe₂O₃H₂O

Fe                                               ⟶       Fe³⁺ + 3e^‑  (Oxidation)

O₂ + 2e^-                                           ⟶       O^2- (Reduction)

Q. Write few lines on Hydrogen peroxide.

INTRODUCTION TO HYDROGEN PEROXIDE

Oxygen combines with hydrogen to form two major hydrides. H₂O and H₂O₂ that is hydrogen peroxide and most common water. The nard was the first person to prepare H₂O₂.

In nature, hydrogen peroxide occur only traces in snow and dew as well as in the air and water which exposed to brilliant sun shine.

 

Q. Describe the laboratory method of preparation of Hydrogen peroxide.

LABORATORY METHOD

H₂O₂ is usually prepared in laboratory by the action of a dilute acid on the peroxide of certain metal.

Acid + Metallic peroxide               ⟶       Salt + H₂O₂

Barium chloride and dilute sulphuric acid H₂SO₄ are used because BaO₂ is insoluble and easily be flittered.

BaO₂ + H₂SO₄            ⟶       BaSO₄ + H₂O₂

 

Q. Describe the industrial method of preparation of Hydrogen peroxide.

INDUSTRIAL PREPRATION

Commercially on large scale H₂O₂ is prepared by the oxidation of isopropyl alcohol with oxygen under reduced pressure.

2C₃H₇OH + O₂           ⟶       2H₂O₂ + 2C₃H₆                                          

 

Q. Describe the physical properties of Hydrogen peroxide.

PHYSICAL PROPERTIES

1.     Pure hydrogen peroxide is a pale blue spray liquid.

2.     It mixes with water to give slightly acidic solution.

3.     Its boiling point is 150^o C but it boils with decompositions.

4.     Its freezing point is about -0.9^o C.

 

Q. Describe the chemical properties of Hydrogen peroxide.

CHEMICAL PROPERTIES

 

1)   Decomposition

When H₂O₂ is exposed to air, it decomposes to form water and oxygen which is exothermic reaction.

2H₂O₂                               ⟶       H₂O + O₂ + Heat

 

2)   As an oxidizing agent

·        Hydrogen peroxide is a strong oxidizing agent because it can readily donate oxygen or accept electron.

H₂O₂                                                   ⟶       H₂O + O                     (Donate Oxygen)

H₂O₂ + 2H + 2e^-                       ⟶       2H₂O                                  (Accept e^-)

·        H₂O₂ liberates iodine from K iodide in the presence of H₂SO₄.

H₂O₂ + 2KI + H₂SO₄             ⟶       K₂SO₄ + 2H₂O + I₂

·        H₂O₂ react with hydrogen supplied to give yellow precipitate of sculpture and water.

H₂O₂ + H₂S                                   ⟶       2H₂O + S

 

3)   As a reducing agent

·        It is also strong oxidizing agent. It can also behave as a reducing agent when it reacts with more powerful oxidizing agent.

·        Hydrogen peroxide reduces chlorine to hydrochloric acid.

Cl₂ + H₂O₂     ⟶       2HCl + O₂

2KMnO₄ + 3H₂SO₄ + 5H₂O₂        ⟶       K₂SO₄ + 2MnSO₄ + 8H₂O + 5O₂

 

Q. Describe the uses of Hydrogen peroxide.

USES OF H₂O₂

·        H₂O₂ is used as mouth washer.

·        It is also used for cleaning wounds as mild antiseptics.

·        It is used for bleaching delicate material such as silk, wool, feathers and human hairs.

·        Liquid H₂O₂ is used for providing oxygen for burning of fuel in space rocket.

·        It is also used in burning of diesel oil.

 

Q. Define Oxidation in three different ways.

OXIDATION

Oxidation can be defined in following three ways.

 

1)   ADDITION OF OXYGEN

Oxidation is a reaction in which oxygen combine with other element or substance to produce oxides.

Rusting Iron

4Fe + 3O₂                      ⟶       2FeO₂

C + O₂                                                ⟶       CO₂

2Mg + O₂                       ⟶       2MgO

2NO + O₂                       ⟶       2NO₂

 

2)   REMOVAL OF HYDROGEN

Oxygen combines with other compounds having hydrogen. It removes hydrogen from them. Removal of H₂ from compound is called oxidation.

H₂S + Cl₂                        ⟶       S + HCl

MnO₂ + 4HCl             ⟶       MnCl₂ + Cl₂ + 2H₂O

 

3)   LOSS OF ELECTRON

Oxidation is also depending on the removal of electron from the substance.

Sn         ⟶       Sn²⁺ + 2e^-

Fe         ⟶       Fe³⁺ + 3e^-

Al          ⟶       Al³⁺ + 3e^-

Sn         ⟶       Sn⁴⁺ + 4e^-

Fe         ⟶       Fe²⁺ + 2e^-

 

Q. Define Reduction in three different ways.

REDUCTION

Reduction is opposite of oxidation reaction. Following are the three definitions of reduction.

 

1)   REMOVAL OF OXYGEN

Reduction means removal of oxygen from metal oxides.

CuO + H₂        ⟶       Cu + H₂O

ZnO + H₂        ⟶       Zn + H₂O

PbO + H₂        ⟶       Pb + H₂O

 

2)   ADDITION OF HYDROGEN

Reduction means addition of hydrogen to an element.

H₂S + Cl₂        ⟶       S + 2HCl

H₂  + Cl₂          ⟶       2HCl

C₂H₄ + H₂      ⟶       C₂H₆

 

3)   GAIN OF ELECTRON

The reaction in which a substance gains electron is called reduction reaction.

Ferric ion reduces to ferrous ion

Fe³⁺ + e^-          ⟶       Fe²⁺

Ferric ion reduces to iron

Fe²⁺ + 2e^-      ⟶       Fe

Zinc reacts with H₂SO₄ to form zinc ion and hydrogen gas

H₂SO₄                            ⟶       2H⁺ + SO₄²⁺

Zn                       ⟶       Zn²⁺ + 2e^-

 

Q. Write note on Ozone.

OZONE

It is a pale blue poisonous gas with a sharp, irritating odor. It is a allotropic form of oxygen with the formula O₃.

 

OCCURENCE

Ozone is formed from atmospheric oxygen by lightning flashes, in nature ozone is very unstable and dissociated readily to form oxygen atom.

O₃         ⟶       O₂ + O      ∆H = -107Kg/mol

Ozone exists in a layer at a height of 20Km above the earth. This ozone layer absorbs too much ultraviolet radiations of the sun and in this way it protects the earth’s surface. Very small amount of ozone is produced around electrical machines during their operation.

 

Q. Write physical properties of Ozone.

PHYSICAL PROPERTIES OF OZONE

1.     Ozone is the pale blue gas at ordinary temperature and pressure.

2.     Ozone has a characteristic smell like that of dilute chlorine.

3.     Ozone is poisonous at conc. above 100 parts per million.

4.     It is only slightly soluble in water but dissolve in turpentine oil.

5.     Pure ozone can be obtained in blue liquid.

 

Q. Write chemical properties of Ozone.

CHEMICAL PROPERTIES

 

1)   Decomposition

Ozone decomposes in ordinary oxygen on heating or even just on standing.

2O₃                     ⟶       3O₂

 

2)   Oxidizing agent

Ozone is more powerful oxidizing agent then oxygen. Oxygen is formed as by production all oxidation reaction.

PbS + 4O₃      ⟶       PbSO₄ + 4O₂

H₂S + 2O₃      ⟶       H₂SO₄ + O₂

SO₂ + O₃          ⟶       SO₃ + O₂

2KI + O₃ + H₂SO₄⟶         K₂SO₄ + I₂ + H₂O + O₂

 

Q. Describe the importance of Ozone.

IMPORTANCE OF OZONE

Ozone is very important from preventing us from ultra violet radiations from the sun. It protect us from the dangerous or harmful effects of such radiations.

 

USES

1.     Ozone is sometimes used for the treatment of domestic water in place of chlorine. It kills bacteria and oxidizes organic compounds present in water.

2.     It is used as bleaching agent because all oxidizing agent are also good bleaching agent.

3.     It is mostly used in preparation of pharmaceuticals, synthetic lubricants. It is also useful commercial organic compound.

 

Q. Describe the laboratory preparation of Ammonia.

LABORATORY PREPRATION

In laboratory it is prepared by heating ammonium salt usually ammonium chloride with slaked lime.

2NH₄Cl + Ca(OH)₂               ⟶       CaCl₂ + 2H₂O + 2NH₃

Both the reactants are solid, so in order to provide the maximum surface area for the reaction, they must be thoroughly grounded.

 

Q. Describe the industrial preparation of Ammonia.

INDUSTRIAL PREPRATION (HABER BOSCH PROCESS)

Ammonia is prepared industrially by Haber Bosch’s process.

PROCESS

1.     In this process a mixture of pure nitrogen and hydrogen in the ratio 1:3 by volumes allowed to react.

2.     The basic problem in NH₃ synthesis is that it is a reversible reaction describe as,
N₂ + 3H₂                ⟶       2NH₃  ∆H = -93kg/mol

3.     The optimum condition of temperature is kept 400^oC—450^oC that of pressure 200—250 atm.

4.     Approximate catalyst ferric oxide (Fe₂O₃) with small amount of Al₂O₃, CaO, K₂O are required to get the maximum yield of ammonia

5.     Ammonia thus obtained in liquefies state after cooling.

6.     The unused gases are recalculated over catalyst for further formation of ammonia.

 

Q. Describe the Physical properties of Ammonia.

PHYSICAL PROPERTIES OF AMMONIA

1.     Ammonia is a colorless gas with a pungent smell.

2.     It is highly soluble in water.

3.     Its solution is alkaline as it turns red litmus paper to blue.

4.     It easily liquefies at ordinary temperature.

5.     Large quantity of ammonia is poisonous and can cause damage to the respiratory system.

 

Q. Describe the Chemical properties of Ammonia.

CHEMICAL PROPERTIES OF AMMONIA

 

1)   Thermal decomposition

Ammonia decomposes into its constituent at high temperature

2NH­₃                                 ⟶       N₂ + 3H₂

 

2)   Reaction with acids

It reacts with acid to give ammonium salts as ammonia is a base.

2NH₃ + H₂SO₄           ⟶       (NH₄)₂SO₄

NH₃ + HCl                     ⟶       NH₄Cl

NH₃ + HNO₃                                ⟶       NH₄NO₃

 

3)   Reaction with CO₂

Ammonia reacts with CO₂ at about 15^oC and about atmosphere to produce urea.

2 NH₃ + CO₂                                ⟶       (NH₂)₂CO + H₂O

 

4)   Reaction with chlorine

If ammonia in excess it first reduce chlorine to produce hydrogen chloride then it react with excess ammonia to produce dense form of ammonium chloride.

2 NH₃ + 3Cl₂              ⟶       6HCl + N₂

6 NH₃ + 6HCl            ⟶       6NH₄Cl

On the other hand if chlorine in excess nitrogen chloride a dangerously explosive and only liquid, it formed.

NH₃ + 3Cl₂                   ⟶       NCl₃ + 3HCl

 

5)   Reaction with O₂

Ammonia does not burn in air, but it burn readily with a greenish yellow flame to form water vapors and nitrogen

4NH₃ + 3O₂                                 ⟶       6H₂O + 2N₂

 

6)   Reaction with water

Ammonia dissolves in water to give ammonium hydroxide solution which is a weak alkali.

NH₃ + H₂O                    ⟶       NH₄OH

 

7)   Reducing agent

Ammonia is not a strong reducing agent. It reduces heated cooper oxide to free copper metal with evolution of nitrogen gas and water

3CuO + 2 NH₃          ⟶       3Cu + N₂ + 3H₂O

 

Q. Describe the uses of Ammonia.

USES OF AMMONIA

1.     Ammonia is used in the manufacture of fertilizers.

2.     It is used in softening of temporary hard water.

3.     Liquid ammonia is used as cooling agent in some refrigerator.

4.     It is used in the manufacture of washing soda by ammonia Solvay process.

5.     It is used as solvent in laundries for removing oil grebes stain from clothes.

 

Q. Write few lines on Nitric acid.

NITRIC ACID

Nitric acid is a very important acid which is used extensively in the laboratory and industry. Nitric acid was first prepared from Sulphuric acid and potassium nitrate by Glibber in 1685.

 

 

 

Q. Describe the laboratory preparation of Nitric acid.

LABORATORY PREPRATION

Nitric acid prepared in the laboratory by heating solid potassium nitrate with concentrated sulphuric acid.

KNO₃ + H₂SO₄           ⟶       KHSO₄ + HNO₃

 

Q. Describe the industry preparation of Nitric acid.

INDUSTRIAL PREPRATION OF NITRIC ACID

OSTWALD’S METHOD

Large amount of nitric acid are made commercially by the Ostwald method. It involves 3 steps.

1.     Ammonia is oxidized by air in the presence of catalyst to form nitric acid. The mixture of ammonia and air is heated to 600^oC with a catalyst platinum.
4 NH₃ + 5O₂       ⟶       4NO + 6H₂O

2.     More air is introduced into the system it oxidizes nitric acid to nitrogen dioxide.
2NO + O₂               ⟶       2NO₂

3.     After cooling NO₂ is passed through water, forming a solution of nitric acid.
3NO₂ + H₂O         ⟶       2HNO₃ + NO

 

Nitric acid obtained from this process is 68% concentrated which can be concentrated further by passing over concentrated H₂SO₄ up to 98%.

 

Q. Describe the physical properties of Nitric acid.

PHYSICAL PROPERTIES

1.     Nitric acid is a colorless fuming liquid with a sharp chocking smell.

2.     It has sour taste.

3.     The density of pure nitric acid is 1.52gm/cm³.

4.     It boils at 83^oC.

5.     It is miscible in water in all proportions.

 

Q. Describe the chemical properties of Nitric acid.

CHEMICAL PROPERTIES

 

1)   Reaction of HNO₃ as acid

Nitric acid is strong mono basic acid and it ionizes completely in water as

HNO₃ + H₂O                ⟶       H₃O⁺ + NO₃^-

2)   Reaction with alkalis

HNO₃ react with alkalis to form salt and water.

HNO₃  + NaOH          ⟶       NaNO₃ + H₂O

HNO₃ + KOH              ⟶       KNO₃ + H₂O

 

3)   Reaction with metal oxides

CaO + 2 HNO₃           ⟶       Ca(NO₃)₂ + H₂O

PbO + 2HNO₃            ⟶       Pb(NO₃)₂ + H₂O

 

4)   Reaction with metal carbonates

CaCO₃ + 2 HNO₃     ⟶       Ca(NO₃)₂ + CO₂ + H₂O

 

5)   Reaction with metals

Mg + 2 HNO₃             ⟶       Mg(NO₃) + H₂

Cu + 4HNO_3(conc)    ⟶       Cu(NO₃)₂ + 2NO₂ + 2H₂O

3Cu + 8HNO_3(dil)    ⟶       3Cu(NO₃)₂ + 2NO₂ + 4H₂O

Pb + 4HNO_3(conc)     ⟶       Pb(NO₃)₂ + 2NO₂ + 2H₂O

3Pb + 8HNO_3(dil)    ⟶       3Pb(NO₃)₂ + 2NO₂ + 4H₂O

Zn + 4HNO_3(conc)     ⟶       Zn(NO₃)₂ + 2NO₂ + 2H₂O

4Zn + 8HNO_3(dil)    ⟶       4Zn(NO₃)₂ + NH₄NO₃ + 4H₂O

 

6)   Reaction with non-metals

Hot concentrated HNO₃ react and it self-redact to NO₂ gas.

C + 4 HNO₃                 ⟶       CO₂ + 4NO₂ + 2H₂O

Si + 4 HNO₃                 ⟶       SiO₂ + 4NO₂ + 2H₂O

P + 5 HNO₃                  ⟶       H₃PO₄ + 5NO₂ + H₂O

 

7)   Reaction with metals

Cu + 4 HNO₃              ⟶       Cu(NO₃)₂ + 2NO₂ + 2H₂O

3Cu + 8 HNO₃           ⟶       3Cu(NO₃)₂ + 2NO + 4H₂O

Pb + 4 HNO₃              ⟶       Pb(NO₃)₂ + 2NO₂ + 2H₂O

 

8)   Reaction with some reducing agents

H₂S + 2 HNO₃            ⟶       S + 2NO₂ + 2H₂O

6FeSO₄ + 2 HNO₃ + 3H₂SO₄        ⟶       3Fe₂(SO₄)₃ + 2NO + 4H₂O

SO₂ + 2 HNO₃            ⟶       H₂SO₄ + 2NO₂

 

9)   As nitrating agent

HNO₃ react with benzene to replace hydrogen atom by the nitro group to form sub situated product into nitro benzene (C₆H₅NO₂)

C₆H₆ + HNO₃              ⟶       C₆H₅NO₂ + H₂O

 

Q. Describe the uses of Nitric acid.

USES OF NITRIC ACID

1.     It is used as an oxidizing agent.

2.     It is used as laboratory reagent.

3.     It is used in manufacture of explosives.

4.     It is used in manufacture dyes.

5.     It is used in manufacture of fertilizers such as ammonium nitrate.

 

Q. Write note on Aqua Regia.

AQUA REGIA

The noble metals like gold and platinum, which are not soluble in concentrated nitric acid, however they dissolve in a mixture of concentrated nitric acid and concentrated HCl taken in the ratio of 13. The mixture is called aqua regain or royal water.

Aqua regain dissolve gold due to liberation of nascent chlorine which form gold chloride with it, which is soluble.

HNO­₃+3HCl                                ⟶       NOCl + 2H₂O + 2Cl           (Nascent chlorine)

NOCl                                  ⟶       NO + Cl                                 (Nitrosyl Chloride)

Au+3Cl                        ⟶       AuCl₃                                                 (Gold chloride)

 

Choose the correct option for each of the following statement.

i.                   When ammonium chloride is heated with base, the gas liberated is:

a.      Ammonia

b.     Nitrogen

c.      Oxygen

d.     Chlorine

ii.                The catalyst used for catalytic oxidation of NH₃ in Ostwald’s method is:

a.      Nickel

b.     Platinum

c.      Chromium

_d.              V₂O₅

iii.             Urea is produced by heating CO₂ with:

a.      Nitric acid

b.     Ammonia

c.      Hydrogen

d.     Potassium

iv.              The most abundant element found in nature is:

a.      Oxygen

b.     Silicon

c.      Nitrogen

d.     Hydrogen

v.                 Hydrogen per oxide is produced in laboratory by heating H₂SO₄ with:

a.      Sodium per oxide

b.     Barium per oxide

c.      Potassium per oxide

d.     Strontium per oxide

vi.              The air we breathe in, usually contains a higher proportion of:

a.      Nitrogen

b.     Oxygen

c.      Carbon dioxide

d.     Water vapors