146

CHAPTER 10
NITROGEN AND ITS COMPOUNDS
10.1 Occurrence
Nitrogen forms the major constituent of the atmosphere, being present to the extent of
about 79 per cent by volume or 76 per cent by weight of the air. In chemical combination
with other elements it occurs in sodium nitrate (Chile saltpetre), in proteins, and many
other animal and vegetable products.

10.2 Methods of Preparation of Nitrogen

Preparation of nitrogen from the atmosphere
The most important gases present in dry air are oxygen (about 21% by volume), carbon
dioxide (about 0.03 % by volume), and atmospheric nitrogen (about 79% by volume).

Fig. 10.1 Preparation of nitrogen from the atmosphere

copper + air copper (II) oxide + nitrogen

2Cu + (N2+O2) 2CuO + N2

sodium + carbon sodium + water

hydroxide + dioxide carbonate
2NaOH + CO2 Na2CO3 + H2O

The apparatus is set up as shown in Fig. 10.1 Water is allowed to flow into bottle
A at a slow rate. Air from this bottle is displaced and made to pass through another bottle
B containing caustic soda solution to remove carbon dioxide.
The carbon dioxide free air is then passed over heated copper which absorbs
oxygen to form copper (II) oxide.
 147

The nitrogen, so freed from oxygen, comes out from the combustion tube. Since
the nitrogen is only slightly soluble in water, it is collected by the downward
displacement of water.
If the nitrogen is required dry. it may be passed through a U-tube containing glass
beads wetted with concentrated sulphuric acid and then collected in a syringe.
Note: The product of this experiment is almost pure nitrogen. It contains about 1% by volume
of the "noble gases ", chiefly argon. The removal of these gases is not possible by
chemical methods. The presence of these gases makes "atmospheric nitrogen" slightly
denser than the pure gas.

Laboratory preparation of nitrogen

Nitrogen gas may be prepared by heating a solution of ammonium nitrite. However,
ammonium nitrite decomposes slowly at ordinary temperatures, so that neither
ammonium nitrite itself, nor its solution in water, should be kept in stock. The
ammonium nitrite is obtained by a reaction between sodium nitrite and ammonium
chloride solutions.

Fig. 10.2 Laboratory preparation of nitrogen

sodium aimmonium ammonium

+ sodium chloride +
nitrite chloride nitrite
NaNO2 + NH4Cl NaCl + NH4NO2

ammonium nitrite nitrogen + water

NH4NO2 N2 + 2H2O

The apparatus is set up as shown in Fig. 10.2

 148

Concentrated solutions of sodium nitrite and ammonium chloride are mixed in a round-
bottomed flask. The flask is then heated gently.
Decomposition to nitrogen occurs as the solution becomes warm. Since nitrogen
gas is slightly soluble in water, it may be collected over water. The nitrogen gas may be
dried as described in the previous experiment.

Other chemical methods of preparation of nitrogen

1. The reaction of chlorine and excess ammonia
chlorine + ammonia nitrogen + ammonium chloride
3Cl2 + 8NH3 N2 + 6NH4Cl
2. Passing ammonia gas over heated copper (II) oxide
ammonia + copper (II) oxide copper + nitrogen + water
2NH3 + 3CuO 3Cu + N2 + 3H2O
3. Reduction of oxides of nitrogen by heated copper
nitrogen
copper + copper (II) oxide + nitrogen
oxide
Cu + 2NO 2CuO + N2

Manufacture of nitrogen
Nitrogen is obtained in industry by the fractional distillation of liquid air. This process is
described in Chapter 10. When liquid air is distilled nitrogen boils off at - 196 C(77K).
Oxygen is left as liquid since its boiling point is higher than that of nitrogen. The
separated nitrogen is then reliquefied.

10.3 Properties of Nitrogen

Physical properties
1. Nitrogen is colourless and odourless. It is slightly less dense than air and only
slightly soluble in water. It condenses to a liquid (b.p. - 196 C), and freezes to a
colourless solid (m.p.-210 C).
2. Nitrogen gas is relatively inert and does not support combustion.

Chemical properties
1. Nitrogen combines with hydrogen and oxygen at higher temperatures.
finely divided powder Fe
nitrogen + hydrogen ammonia
200atm,450°C

N2 + 3H2 2NH3

nitrogen + oxygen nitrogen oxide

N2 + O2 2NO
 149

2. It combines with many metals, on heating to a dull red heat or higher

temperatures, to form nitrides.

magnesium + nitrogen magnesium nitride

3Mg + N2 Mg3N2

Nitrides react with water to liberate the ammonia gas.

magnesium magnesium
+ water ammonia +
nitride hydroxide
Mg3 N2 + 6H2O 2NH3 + 3Mg(OH)2

3. It combines directly with calcium carbide at about l000 C to form calcium

cyanamide, CaCN2, an important nitrogenous fertilizer.

calcium
+ nitrogen calcium cyanamide + carbon
carbide
1000 C
CaC2 + N2 CaCN2 + C

A mixture of calcium cyanamide and carbon is known as "nitrolime". Calcium

cyanamide reacts with water to form ammonia.

calcium
+ water calcium carbonate + ammonia
cyanamide
CaCN2 + 3H2O CaCO3 + 2NH3

10.4 Uses of Nitrogen

This principal use of nitrogen is in the manufacture of ammonia, from which nitrogenous
fertilizers, nitric acid and urea are now mainly prepared.

10.5 Compounds of Nitrogen

Ammonia NH3
This hydride of nitrogen, NH3, can be formed by decomposition of nitrogenous materials.
This decomposition is brought about by bacteria.

Preparation of ammonia
Laboratory preparation of ammonia
Ammonia may be prepared in the laboratory by heating any ammonium salt with an
alkali. Usually, a mixture of ammonium chloride and calcium hydroxide (slaked lime) is
used.
 150

.
Fig. 10.3 Laboratory preparation of ammonia
calcium ammonium calcium
+ + water + ammonia
hydroxide chloride chloride
Ca(OH)2 + 2NH4Cl CaCl2 + 2H2O 2NH3

An excess of the slaked lime and ammonium chloride are mixed. The mixture is ground
by using a mortar and pestle, and placed in a round-bottomed flask. The apparatus is set
up as shown in Fig. 10.3.When this mixture is heated ammonia gas is evolved.
It is dried by passage through a lime tower which contains the drying agent,
quicklime, CaO.
Since ammonia is lighter than air and very soluble in water, it is collected by
downward displacement of air. The gas jar is known to be filled with ammonia if a moist
red litmus paper placed at the mouth of the gas jar turns blue.
Note: Ammonia gas should not be dried by usual drying agents, such as, calcium chloride,
concentrated sulphuric acid and phosphorus (V) oxide. These compounds react with
ammonia.

Manufacture of ammonia
Haber Process
Ammonia is obtained on an industrial scale by the Haber Process. This process is
based on the direct combination of nitrogen and hydrogen.
nitrogen + hydrogen ammonia
N2 + 3H2 2NH3

This is a reversible reaction.

For the process to be economically successful, the yield of the ammonia should be
increased by driving the reaction towards the right.
This can be done
(a) by conducting the reaction at the lowest possible temperature,
 151

(b) by increasing the pressure and

(c) by using a suitable catalyst to reduce the time required for the reaction.

Fig. 10.4 Synthesis of ammonia

The mixture of nitrogen and hydrogen, in the ratio of one to three by volume, is
compressed to 200 atmospheric pressure and passed over the catalyst heated at 450
C( 723 K). The catalyst used is usually finely divided reduced iron impregnated with
alumina, Al2O3 . The hot gases leaving the catalyst chamber are passed through a heat
exchanger in order to preheat the incoming hydrogen - nitrogen mixture. After, washing
out the ammonia, or comperssing and removing it as the liquid, the residual gases are
mixed with more hydrogen and nitrogen and again passed through the catalyst Chamber.
The recirculation is repeated as required.
The ammonia produced is liquefied and stored for further use.

Test for ammonia

Ammonia gas has a characteristic pungent smell. It turns moist red litmus paper blue.
Dense white fumes are formed when the gas comes into contact with hydrogen chloride
gas or concentrated hydrochloric acid.

Physical properties of ammonia

1. Ammonia is a colourless gas with a pungent odour. It is lighter than air.
2. It is extremely soluble in water. This considerable solubility is often demonstrated in
the laboratory by the "fountain experiment" as illustrated in Fig. 10.5
The glass vessel is filled with dry ammonia gas and the apparatus arranged as Shown.
Some of the air inside the vessel is expelled by gently warming the bulb. On allowing the
vessel to cool, a little water is drawn up the capillary tube and is discharged into the flask.
This water immediately dissolves practically the whole of the ammonia gas in the flask to
 152

create a partial vacuum. Water rushes up the capillary the flask to fill the vacuum and a
fountain like effect is produced.

.
Fig. 10.5 Fountain experiment
The alkaline nature of ammonia can be shown in this experiment by adding a red litmus
solution to the water in the trough. When the ammonia in the glass vessel dissolves in
the litmus solution, it is turned blue.

Chemical properties of ammonia

1. Ammonia gas reacts with hydrogen chloride gas to form solid ammonium chloride.
hydrogen ammonium chloride
ammonia +
chloride (dense fumes)
NH3 + HCl NH4Cl
2. Dry ammonia will burn in oxygen or in an atmosphere of air slightly enriched by
oxygen but not in air alone.
ammonia + oxygen nitrogen + water
4NH3 + 3O2 2N2 + 6H2O
3. When ammonia is passed over heated copper (II) oxide, it reduces the copper (II)
oxide to copper.
copper (II) oxdide + ammonia copper + water + nitrogen
3CuO + 2NH3 3Cu + 3H2O + N2
4. Nitrogen is obtained when excess of ammonia reacts with chlorine,
ammonia + chlorine nitrogen + ammonium chloride
8NH3 + 3C12 N2 + 6NH4Cl
but if the chlorine is in excess, the explosive oily substance, nitrogen trichloride is
formed.
 153

nitrogen hydrogen
ammonia + chlorine +
trichloride chloride
NH3 + 3Cl2 NCl3 + 3HCl
5. Ammonia combines directly with some metals on heating. Thus, dry ammonia
passed over sodium metal heated to red heat forms sodamide. Magnesium however forms
the nitride.
ammonia + sodium sodamide + hydrogen
2NH3 + 2Na 2NaNH2 + H2

ammonia + magnesium magnesium nitride + hydrogen

2NH3 + 3Mg Mg3N2 + 3H2
Both sodamide and magnesium nitride are decomposed by water to regenerate ammonia
gas.
sodamide + water Sodium hydroxide + ammonia
NaNH2 + H2O NaOH + NH3

magnesium magnesium
+ water + ammonia
nitride hydroxide
Mg3N2 + 6H2O 3Mg(OH)2 + 2NH3
6. An ammonium hydroxide solution, obtained by dissolving ammonia gas in water is
also known as ammonia solution.
ammonia + water ammonium hydroxide
NH3 + H2O NH4OH

Uses of ammonia
The liquid of gas is used in refrigerators, and in the large-scale manufacture of
(1) fertilizers, such as ammonium sulphate, ammonium nitrate and ammonium,
phosphate,
(2) nitric acid and nitrates,
(3) urea and
(4) certain other organic compounds.
 154

Composition of ammonia

Fig. 10.6 Eudiometer for dissociation of ammonia

Some dry ammonia gas is enclosed over mercury in a eudiometer tube and the
volume measured at atmospheric pressure (by making levels A and B equal). The original
volume in one such experiment was 20cm 3. On passing electric sparks across the
electrodes, the ammonia decomposes into nitrogen and. hydrogen. After completion of
the decomposition (this is known when no further increase in volume occurs) the electric
current is switched off and the tube is allowed to cool down to room temperature. The
volume is measured at atmospheric pressure by making the two levels of mercury equal.
It is found to be double the original ammonia volume (40cm 3). This increase in volume is
due to the decomposition of ammonia into nitrogen and hydrogen.
A known excess of oxygen (20 cm 3) is then introduced into the eudiometer and
the mixture sparked. (Formation of water results.) After cooling to room temperature, the
volume of residual gases is measured (15cm3). This contraction in volume is due to the
combination of the hydrogen and oxygen to form water. This water is condensed to a
negligible volume. The gaseous residue consists only of excess oxygen and nitrogen.
hydrogen + oxygen water
2H2 + O2 2H2O
2 vol 1 vol
A drop or two of pyrogallol is introduced into the eudiometer. This substance
removes the oxygen by absorption. Nitrogen is left behind and its volume may be
measured (10cm3).
The data obtained may be treated as follows:
Volume of ammonia = 20cm3
Volume of N2+ H2 after sparking = 40cm3
Volume of N2+ H2+O2 added = 60cm3
Volume of O2 = 20cm3
 155

Volume after sparking with excess O2

and cooling or (N2 + excess O2) = 15cm3
Volume after absorption with pyrogallol
(i.e., volume of N2) = l0cm3
Volume of O2 excess = 5cm3
Volume of O2 reacted = (20-5) = 15cm3
Since 1 volume of oxygen reacts with 2 volumes of hydrogen to form water,
Volume of H2 is (2x vol of O2)=30cm3
That is also,
Volume of H2 from ammonia = 30 cm3

nitrogen + hydrogen = ammonia

10 cm3 + 30 cm3 = 20 cm3
1 vol + 3 vol = 2 vol
By applying Avogadro's Theory,
1 molecule + 3molecules = 2 molecules
i.e., 2 atoms + 6 atoms = 2 molecules
l atom of 3 atoms of 1 molecule of
+ =
nitrogen hydrogen ammonia
Formula of ammonia = NH3
Molecular mass of NH3 = 14+3x1 = 17

This is in agreement with the relative vapour density of ammonia which is 8.5.
Molecular mass = 2 x vapour density
= 2 x 8.5
= 17

Preparation of solution of ammonium hydroxide

Fig. 10.7 Preparation of solution of ammonium hydroxide

 156

Some ammonium chloride and calcium hydroxide are placed in a round-bottomed flask.
The apparatus is set up as shown in Fig. 10.7. The rim of the inverted funnel attached to
the end of the delivery tube should just touch the surface of the water in the beaker.
The flask is gently heated. The liberated ammonia gas will dissolve in the water
contained in the beaker. After some time the water in the beaker will be found to have
acquired the smell of ammonia gas which has dissolved in it. The solution is known as
ammonia solution.

Properties of ammonia solution

1. Since ammonia solution has an alkaline property, it will precipitate insoluble
metallic hydroxides when mixed with solutions of salts of me heavy metals.
ammonium iron(III) iron(III) hydroxide ammonium
+ +
hydroxide chloride (brown) chloride
3NH4OH + FeCl3 Fe (OH)3 + 3NH4Cl
Zinc and copper (II) hydroxides will dissolve in excess ammonia solution to give
solutions of complex ammines. The complex ammine obtained from copper has a deep
blue colour.
2. Ammonia solution will neutralize acids to form ammonium salts.
ammonium hydrochloric
+ ammonium chloride + water
hydroxide acid
NH4OH + HCl NH4Cl + H2O

Ammonium salts
Preparation
Neutralization of an ammonia solution which has basic property by an acid yields
ammonium salt and water.

Properties
Ammonium salts of acids having a high proportion of oxygen are usually decomposed by
heat.
ammonium nitrite nitrogen + water
NH4NO2 N2 + 2H2O

ammonium nitrate dinitrogen oxide + water

NH4NO3 N2O + 2H2O
Ammonium chloride is one of the few substances which sublimes, instead of
decomposing, on heating.

Oxides of nitrogen
 157

There are three common oxides of nitrogen. They are dinitrogen oxide or nitrous oxide,
N2O; nitrogen oxide or nitric oxide, NO; and nitrogen dioxide, NO 2 or dinitrogen
tetroxide, N2O4, depending on the conditions.

Dinitrogen oxide N2O

Preparation of dinitrogen oxide
Dinitrogen oxide gas may be prepared by heating ammonium nitrate (or any mixture of
salts which will yield ammonium nitrate on double decomposition).

Fig. 10.8 Laboratory perparation of dinittogen oxide

ammonium potassium ammonium potassium

+ +
sulphate nitrate nitrate sulphate
(NH4)2SO4 + 2KNO3 2NH4NO3 + K2SO4
ammonium nitrate dinitrogen oxide + water
NH4NO3 N2O + 2H2O
A mixture of potassium nitrate and ammonium sulphate is placed in a flask. The
apparatus is set up as shown in Fig. 10.8.
The flask is heated gently. On heating, the ammonium nitrate melts and
effervesces (i.e., gives off bubbles of gas).
The gas liberated is collected over hot water. This gas is fairly soluble in cold
water.
Note: Ammonium nitrate can give dinitrogen oxide on heating. It is, however, likely to explode on
heating. But the reaction is quite safe if ammonium nitrate is generated in the flask by double
decomposition during, the preparation as above.

Test for dinitrogen oxide

Dinitrogen oxide gas rekindles a brightly glowing splint.
Since the above test is similar to the test for oxygen, the two gases cannot be
distinguished by this test. To distinguish them the following test may be carried out.
 158

Invert a jar of the gas over cold water in a trough and shake it. If the gas is
dinitrogen oxide the water level in the gas jar will rise above the level in the trough,
showing the gas to be fairly soluble in water. Oxygen is almost insoluble in water and no
rise in the water level would be observed.
Dinitrogen oxide does not give brown fumes with nitrogen oxide while oxygen
with nitrogen oxide forms brown fumes of nitrogen dioxide.

Physical properties of dinitrogen oxide

1. Dinitrogen oxide is a colourless gas with a sweetish odour. It is fairly soluble in
cold water. It is a neutral oxide.
2. It will relight a glowing splint, and will support combustion. However, if a splint
is feebly glowing it will be extinguished.
To be rekindled, the glowing portion of the splint must be not enough to decompose
some dinitrogen oxide into nitrogen and oxygen. The mixture will then be rich enough in
oxygen to stimulate the combustion of the splint.
dinitrogen oxide nitrogen + oxygen
2N2O 2N2 + O2

Chemical properties of dinitrogen oxide

1. Dinitrogen oxide is readily decomposed above 600 C.

dinitrogen oxide 600 C nitrogen + oxygen

2N2O 2N2 + O2
2. Dinitrogen oxide will support the combustion of those burning materials the
flames of which are not enough to decompose it and so liberate free oxygen with which
the material may combine.
dinitrogen
carbon + carbon dioxdie + nitrogen
oxide
C + 2N2O CO2 + 2N2

dinitrogen phosphorus (v)

phosphorus + + nitrogen
oxide oxide
P4 + 10N2O P4O10 + 10N2
dinitrogen
sodium + sodium oxdie + nitrogen
oxide
2Na + N2O Na2O + N2

dinitrogen magnesium
magnesium + + nitrogen
oxide oxide
Mg + N2O MgO + N2
 159

Uses of dinitrogen oxide

It is used as an anaesthetic for minor surgical operations.
Hysteria sometimes follows recovery from the anaesthetic, and this gave rise to the name
" Laughing gas".

Nitrogen oxide NO
Preparation of nitrogen oxide

Fig. 10.9 Laboratory preparation of nitrogen oxide

nitric acid copper(II) nitrogen

copper + + + water
(dil) nitrate oxide
3Cu + 8HNO3 3Cu(NO3)2 + 2NO + 4H2O
Some copper turnings are placed in a flat-bottomed flask. The apparatus is set up
as shown in Fig. 10.9.
Dilute nitric acid, made by adding an equal volume of concentrated nitric acid to
water, is slowly added to the copper turnings. Vigorous effervescence occurs and the
flask is filled with brown fumes. These brown fumes are nitrogen dioxide, produced
partly by the action of the acid upon the copper and partly by the oxidation of the
nitrogen oxide by the oxygen of the air in the flask.
nitrogen oxide + oxygen nitrogen dioxide
2NO + O2 2NO2
The brown fumes dissolve in the water and the nitrogen oxide is collected as a
colourless gas by the downward displacement of water.
Tests for nitrogen oxide
1. Remove the cover from a gas jar of nitrogen oxide. Reddish brown fumes are
produced immediately when the nitrogen oxide comes into contact with oxygen of the
air.
 160

nitrogen oxide + oxygen nitrogen dioxide

2NO + O2 2NO2

2. Pour a cold acidified solution of iron (II) sulphate into a gas jar of nitrogen oxide.
A dark brown colouration caused by formation of a black compound, nitroso iron (II)
sulphate, FeSO4. NO, is obtained.
nitrogen nitroso iron (II)
iron (II) sulphate +
oxide sulphate
FeSO4 + NO FeSO4.NO

Physical properties of nitrogen oxide

1. It is a colourless gas.
2. It is insoluble in water.
3. It is neutral to litmus.

Chemical properties of nitrogen oxide

1. Nitrogen oxide combines directly with oxygen to form the brown gas, nitrogen
dioxide.
nitrogen oxdie + oxygen nitrogen dioxde
2NO + O2 2NO2
2. Nitrogen oxide is soluble in solutions of iron (II) salts with which it forms brown
nitroso complexes.
nitrogen nitroso iron (II)
iron(II) sulphate +
oxide sulphate
FeSO4 + NO FeSO4NO
3. It combines directly with chlorine in the presence of charcoal to form nitrosyl
chloride, NOCl.
nitrogen oxide + chlorine nitorsyl chloride
2NO + Cl2 2NOCl
4. Nitrogen oxide will support the combustion of those burning materials whose
temperatures are high enough to decompose it and so liberate free oxygen with which the
material may combine.
nitrogen phosphorus (V)
phosphorus + + nitrogen
oxide oxide
P4 + 10NO P4O10 + 5N2

nitrogen magnesium
magnesium + + nitrogen
oxide oxide
2Mg + 2NO 2MgO + N2

Composition of nitrogen oxide

 161

Fig. 10.10 Determination of the composition of nitrogen oxide

A suitable volume of nitrogen oxide is introduced into an apparatus (Fig- 10.10)
made of hard glass tube and its exact volume measured at atmospheric pressure (after
making levels A and B equal). The iron wire is then electrically heated. The measured
volume of nitrogen oxide is decomposed. The iron metal then combines with the oxygen
from nitrogen oxide. Nitrogen from nitrogen oxide is liberated.
nitrogen magnetic iron
iron + + nitrogen
oxide oxide
3Fe + 4NO Fe3O4 + 2N2
When the decomposition is complete there will be no further change in volume.
The electric current is switched off and the tube is allowed to cool. Water level A rises
towards C. The volume of nitrogen at atmospheric pressure is measured by making levels
B and C equal. The volume of nitrogen is found to be one-half of the original volume of
nitrogen oxide, i.e., at constant temperature and pressure.
2 vol of nitrogen oxide 1 vol of nitrogen
1 molecule of nitrogen is obtained from 2 molecules of nitrogen oxide.
By applying Avogadro's Theory,
2 atoms of nitrogen ........ 2molecules of nitrogen oxide
1 atom of nitrogen .......... 1 molecule of nitrogen oxide
The formula of nitrogen oxide = NOx
The vapour density of nitrogen oxide = 15
The relative moluecular mass of nitrogen oxide = 15 x 2 = 30
NOx = 30
 162

(14+16)x = 30
x = 1
The molecular formula of nitrogen oxide is NO.

Nitrogen dioxide NO2

Laboratory preparation of nitrogen dioxide

Fig. 10.11 Laboratory preparation of nitrogen dioxide

nitric acid copper (II) nitrogen

copper + + water +
(cone) nitrate dioxide
Cu + 4HNO3 Cu(NO3)2 + 2H2O + 2NO2

Some copper turnings are placed in a round-bottomed flask. The apparatus is set
up as shown in Fig. 10.11.
Some concentrated nitric acid is cautiously added. The flask is heated gently.
Since nitrogen dioxide reacts with water and is heavier than air, the gas is collected
by the upward displacement of air.
Since nitrogen dioxide is reddish brown, it is easily observed when the jar is full.
 163

Preparation of nitrogen dioxide from lead (II) nitrate

Fig. 10.12 Laboratory preparation of nitrogen dioxide from lead II nitrate

lead (II) lead (II) nitrogen
+ + oxygen
nitrate oxide dioxide
2Pb(NO3)2 2PbO + 4NO2 + O2
Some lead (II) nitrate is placed in a hard glass tube. The apparatus is set up as
shown in Fig.l0.l2.
On heating, the lead (II) nitrate decomposes to give off nitrogen dioxide which is
a brown gas. This gas is liquefied in a U-tube immersed in a freezing mixture of ice and
salt.

Physical properties of nitrogen dioxide

1. Nitrogen dioxide is usually seen as a reddish brown gas at room temperature. It may be liquefied at 22 C under normal pressure and solidified at —10 C under the same pressure. The solid is colourless when pure, and the liquid has a pale yellow colour. The

depth of colour increases as the temperature is increased from the melting point to about 150 C. These colour changes result from the following dissociation reactions.

heat
dinitrogen tetroxide nitrogen dioxdie
cool
N2O4 2NO2
(light yellow) (dark brown)
This reversible reaction is an example of thermal dissociation.
2. It has a pungent, irritating smell.

Chemical properties of nitrogen dioxide

1. Thermal dissociation of N2O4
At —10 C, nitrogen dioxide exists as a solid. The solid consists of the colourless "double"
molecules N2O4, dinitrogen tetroxide. These molecules dissociate on heating to form the
brown NO2 molecules. At above 150 C it begins to dissociate into the colourless nitrogen
oxide and oxygen. The dissociation is complete at the temperature of 620 C.
dinitrogen nitorgen
nitrogen oxide + oxygen
tetroxide dioxide
N2O4 2NO2 2NO + O2
 164

Since all the above reactions or dissociations are reversible, reducing the temperature will
cause the reverse reactions to occur.
2. It dissolves in water to form a mixture of nitric and nitrous acids, and may
therefore be regarded as a mixed anhydride.
nitrogen
+ water nitric acid + nitrous acid
dioxide
2NO2 + H2O HNO3 + HNO2
3. When nitrogen dioxide is absorbed in aqueous alkalis, the corresponding nitrates
and nitrites are formed.
sodium nitrogen sodium sodium
+ + + water
hydroxide dioxide nitrate nitrite
2NaOH + 2NO2 NaNO3 + NaNO2 + H2O

Laboratory preparation of nitric acid, HNO3

Fig, 10.13 Laboratory preparation of nitric acid

potassium sulphuric potassium

+ + nitric acid
nitrate acid(conc) hydrogensulphate
KNO3 + H2SO4 KHSO4 + HNO3

Some potassium nitrate crystals are placed in the bulb of the retort. Concentrated
sulphuric acid is added just to cover the potassium nitrate crystals.
The retort is gently heated. Effervescence will occur with the evolution of brown
fumes. The brown fumes are nitrogen dioxide formed by the slight decomposition of the
nitric acid by heat.

nitric acid nitrogen

water + + oxygen
(cone) dioxide
4HNO3 2H2O + 4NO2 + O2
 165

The nitric acid is collected as a yellow liquid in the receiver cooled under running
water.
In this reaction the less volatile sulphuric acid has displaced the more volatile
nitric acid from its salt.

Manufacture of nitric acid

Nitric acid is manufactured by the Ammonia-Oxidation Process.

Fig. 10.14 Catalytic oxidation of ammonia in the manufacture of nitric acid

The oxidation of ammonia is carried out by passing ammonia gas and excess of air
through a multilayered very fine gauze consisting of platinum (90%) rhodium (10%)
catalyst. The reaction is initiated by heating the catalyst to red heat. Once the reaction is
started the heat generated maintains the reaction temperature.
Pt-Rh
ammonia + oxygen nitrogen oxide + water
red heat
4NH3 + 5O2 4NO + 6H2O
The nitrogen oxide so formed is rapidly cooled and made to combine with oxygen
from excess of air to form nitrogen dioxide.
nitrogen oxide + oxygen nitrogen dioxide
2NO + O2 2NO2
The nitrogen dioxide in the presence of more air, is then absorbed in hot water to
yield nitric acid.
nitrogen
water + + oxygen nitric acid
dioxide
2H2O + 4NO2 + O2 4HNO3
 166

Physical properties of nitric acid

Nitric acid is a colourless, fuming liquid.

Chemical properties of nitric acid

It is a very strong acid as well as a powerful oxidizing agent.
(a) Nitric acid acting as an acid
(i) It neutralizes bases, forming metallic nitrates.
potassium nitric acid potassium
+ + water
hydroxide (dil) nitrate
KOH + HNO3 KNO3 + H2O
copper (II) nitric copper (II)
+ + water
oxide acid(dil) nitrate
CuO + 2HNO3 Cu (NO3)2 + H2O
(ii) It liberates carbon dioxide in reaction with metallic carbonates and
hydrogencarbonates.
calcium nitric calcium carbon
+ + + water
carbonate acid(dil) nitrate dioxide
CaCO3 + 2HNO3 Ca(NO3)2 + CO2 + H2O
sodium hydro- nitric sodium carbon
+ + + water
gencarbonate acid(dil) nitrate dioxide
NaHCO3 + HNO3 NaNO3 + CO2 + H2O

(b) Nitric acid as an oxidizing agent

(i) It oxidizes metals to give nitrogen dioxide or nitrogen oxide.
Copper + nitric acid copper(II) + water + nitrogen
conc nitrate dioxide
Cu + 4 HNO3 Cu(NO3)2 + 2 H2O + 2 NO2
nitric acid copper(II) nitrogen
+ + water +
.copper (dil) nitrate oxide
3Cu + 8HNO3 3Cu(NO3)2 + 4H2O + 2NO
Other metals react similarly to give oxides of nitrogen, the nature of the gaseous
product depending on the metal, the concentration of the acid and the temperature
employed. (See the following section.)
(ii) Nitric acid can also oxidize, certain non-metallic element and certain compounds.
iron (II) sulphuric nitric iron(III) nitrogen
+ + + water +
sulphate acid acid sulphate oxide
6FeSO4 + 3H2SO4 + 2HNO3 3Fe2(SO4)3 + 4H2O + 2NO
In the given example, the nitric acid has oxidized the green iron (II) sulphate to
brown iron (III) sulphate.
 167

Further examples, of the oxidizing actions of hot, concentrated nitric acid are the
following.
nitric acid nitrogen carbon
carbon + water + +
(conc) dioxide dioxide
C + 4HNO3 2H2O + 4NO2 + CO2
nitric phosphoric nitrogen
phosphorus + + + water
acid(conc) acid dioxide
P4 + 20HNO3 4H3PO4 + 20NO2 + 4H2O
nitric acid sulphuric nitrogen
sulphur + + + water
(conc) acid dioxide
S + 6HNO3 H2SO4 + 6NO2 + 2H2O
nitric acid nitrogen
iodine + iodic acid + + water
(conc) dioxide
I2 + 10HNO3 2HIO3 + 10NO2 + 4H2O
Hydrogen nitric acid nitrogen
+ sulphur + + water
sulphide (dil) oxide
3H2S + 2HNO3 3S + 2NO + 4H2O
Hydrogen nitric acid nitrogen
+ iodine + + water
iodide (dil) oxide
6HI + 2HNO3 3I2 + 2NO + 4H2O
(c) Action of nitric acid on the metals
The action of nitric acid on the metals is particularly interesting in view of the
number of different products that may be obtained. The only metals not attacked by the
acid are the noble metals like gold and platinum, although certain other metals, such as
iron and chromium become passive when treated with the strong acid. This passivity is
thought to be due to the formation of an insoluble oxide film, which prevents further
attack by the acid.
All the other metals react with nitric acid and the products of the reaction appear
to be dependent on the reactivity of the metal, as determined by its position in the activity
series and by the concentration of the nitric acid.
With the metals above hydrogen in the activity series, the primary reaction appears to be the liberation of nascent hydrogen which then reduces the nitric acid. The extent to which the nitric acid is reduced, is determined by the rate of formation of
the nascent hydrogen and by the concentration of nitric acid as indicated by the following equations.

nascent nitrogen
nitric acid + water +
hydrogen dioxide
HNO3 + (H) H2O + NO2
nascent nitrogen
nitric acid + water +
hydrogen oxide
HNO3 + 3(H) 2H2O + NO
nascent dinitrogen
nitric acid + water +
hydrogen oxide
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2HNO3 + 8(H) 5H2O + N2O

nascent
nitric acid + water + hydroxylamine
hydrogen
HNO3 + 6(H) 2H2O + NH2OH

nascent
nitric acid + water + ammonia
hydrogen
HNO3 + 8(H) 3H2O + NH3
(i) With the very reactive magnesium metal and very dilute nitric acid, the rate of
formation of nascent hydrogen is so very excessive that part of it is liberated as hydrogen
gas.

nitric acid magnesium

magnesium + + Hydrogen
(dil) nitrate
Mg + 2HNO3 Mg(NO3)2 + H2

With slightly more concentrated acid the balance is restored, and the main product
is now ammonia, which is however neutralized by the excess acid to form ammonium
nitrate.

magnesium
magnesium + nitric acid (cone) + ammonia + water
nitrate
4Mg + 9HNO3 4Mg(NO3)2 + NH3 + 3H2O
ammonia + nitric acid (cone) ammonium nitrate
NH3 + HNO2 NH4NO3
(ii) With a slight less reactive metal such as zinc or aluminium and hot dilute nitric
acid, no free hydrogen escapes. The main product is ammonium nitrate, but this-may
decompose at the temperature of the reaction to liberate dinitrogen oxide gas.

nitric acid ammonium

zinc + zinc nitrate + + water
(dil) nitrate
4Zn + 10HNO3 4Zn(NO3)2 + NH4NO3 + 3H2O

ammonium nitrate dinitrogen oxide + water

NH4NO3 N2O + 2H2O

(iii) With the metals, such as copper, silver and mercury which are below hydrogen in
the electrochemical series, the reaction is an oxidation reaction. Copper with cold dilute
nitric acid gives mainly nitrogen oxide,
nitric copper (II) nitrogen
copper + + + water
acid(dil) nitrate oxide
3Cu + 8HNO3 3Cu(NO3)2 + 2NO + 4H2O
but with the concentrated acid nitrogen dioxide is the main product.
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nitric acid copper(II) nitrogen

copper + + + water
(conc) nitrate dioxide
Cu + 4HNO3 Cu(NO3)2 + 2NO2 + 2H2O

(iv) In the case of tin with cold dilute nitric acid, the reaction is
nitric acid ammonium
tin + tin(II) nitrate + + water
(dil) nitrate
4Sn + 10HNO3 4Sn(NO3)2 + NH4NO3 + 3H2O

with hot concentrated nitric acid, metastannic acid is produced.

nitric acid metastannic nitrogen
tin + + + water
(conc) acid dioxide
Sn + 4HNO3 H2SnO3 + 4NO2 + H2O

(d) The ability of aqua-regia (3 volumes concentrated HC1 +1 volume concentrated

HNO3) to dissolve the noble metal gold, is due to the oxidation of the hydrochloric acid
by the nitric acid to liberate chlorine which converts the gold into gold (III) chloride. The
latter dissolves in excess of hydrochloric acid to form the soluble chloroauric acid,
HAuCl4.

hydrochloric nitrosyl
+ nitric acid chlorine + + water
acid chloride
3HCl + HNO3 Cl2 + NOCl + 2H2O

goal + chlorine gold (III) chloride

2Au + 3Cl2 2AuCl3

gold (III) chloride + hydrochloric acid chloroauric acid

AuCl3 + HCl HAuCl4

Uses of nitric acid

Nitric acid is used mainly of the manufacture of explosives and dyes. A valuable fertilizer
can be obtained by neutralizing the acid with lime. The calcium nitrate obtained, mixed
with excess lime forms a non-deliquescent basic salt which is applied to the soil.

Nitrates
The properties of the nitrates vary according to the position of the metal in the activity
series.
K Nitrates of these metals are decomposed by heat to All nitrates are
Na the nitrite and oxygen. soluble in water.

e.g., 2KNO3 2KNO2 + O2

Ca
Mg
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Al
Zn
Fe
Nitrates of these metals are-decomposed on heating
Pb
to the oxide of the metal, nitrogen dioxide and
Cu
oxygen.
Hg Nitrates of these metals are decomposed on healing
to the metal, nitrogen dioxide and oxygen.
Ag e.g., Hg(NO3)2 Hg + 2NO2 + O2
Note: Ammonium nitrate is decomposed into dinitrogen oxide and steam on heating.
ammonium nitrate dinitrogen oxide + water
NH4NO3 N2O + 2H2O

Test for nitrate radical

Add a little dilute sulphuric acid and two or three crystals of iron (II) sulphate into a
solution of nitrate salt. Shake to dissolve them. Hold the test tube in a slanting position
and pour a slow continuous stream of concentrated sulphuric acid down the side. It will
form a separate layer underneath the aqueous layer and, at the junction of the two, a
brown ring will be seen.
The formation of this brown ring is the characteristic test for a soluble nitrate and
is known as the Brown Ring Test.
Explanation: The concentrated sulphuric acid and the nitrate react to yield nitric acid.
potassium + sulphuric potassium + nitric
nitrate acid (conc) hydrogensulphate acid
KNO3 + H2SO4 KHSO4 + HNO3

The nitric acid is then reduced by some of the iron (II) sulphate to nitrogen oxide.
iron(II) + nitric + sulphuric iron(III) + water + nitrogen
sulphate acid acid sulphate oxide
6FeSO4 + 2HNO3 + 3H2SO4 3Fe2(SO4)3 + 4H2O + 2NO

The nitrogen oxide reacts with iron (II) sulphate to give the brown compound
FeSO4.NO, which appears as the brown ring.
iron (II) sulphate + nitrogen oxide nitroso iron (II) sulphate
FeSO4 + NO FeSO4.NO

10.6 The Nitrogen Cycle

Nitrogen is an essential constituent of vegetable and animal tissues. Plants absorb nitrates
from the soil. However, leguminous plants (e.g., beans, peas) are able to utilize nitrogen
from the atmosphere to form proteins, which are the nitrogen containing organic
compounds. When a plant dies, compounds of ammonia are produced. These compounds
are attacked by "nitrifying bacteria" in the soil and are there by converted into nitrates
 171

which are retained in the soil. In this way, the cycle is complete.The conversion of
atmospheric nitrogen to nitrates of the soil is termed “Nitrogen-fixing”
Animals, on the other hand, can also make use of nitrogen. The plants are eaten
by animals, and the protein is digested by the animals. An excess of protein is excreted
and is returned to the soil as urea. Bacteria decomposes the urea to ammonia. Decaying
living creatures also yield compounds of ammonia. These ammonia compounds are
attacked by nitrifying bacteria in the soil, and converted into nitrates.
However, because of the denitrifying bacteria, which is also present in the soil,
there is a constant loss of nitrogen from the nitrates in the soil to the atmosphere. The
natural processes by which nitrogen in the atmosphere is returned to the soil maybe
summarized as follows:
Electrical discharges in the atmosphere, such as lightning cause some formation
of nitrogen dioxide.
nitrogen + oxygen nitrogen oxide
N2 + O2 2NO

nitrogen oxide + oxygen nitogen dioxide

2NO + O2 2NO2
The nitrogen dioxide combines with rain and oxygen to give nitric acid which can
react with substances in the soil, such as chalk to form nitrates.
oxygen + water + nitrogen nitric acid
dioxide
O2 + 2H2O + 4NO3 4HNO3

calcium nitric calcium carbon

+ + water +
carbonate acid(dil) nitrate dioxde
CaCO3 + 2HNO3 Ca(NO3)2 + H2O + CO2
The nitrogen cycle is summarized diagrammatically in Fig. 10.15,
Haber process
NH3 electric discharge
atmospheri HNO3
c N2 lighting and rain

peas beans

ammonium Plant nitrifying nitrates

bacteri
salts proteins a in the
soil
 172

ure excretion animal

a proteins
Fig. 10.15 The nitrogen cycle

SUMMARY
Nitrogen is colorless and odourless. It is slightly less dense than air and only
slightly soluble in water. It condenses to a liquid (b.p. - 196 C), and freezes to a
colourless solid (m.p.-210 C).
Nitrogen forms the major constituent of the atmosphere, being present to the
extent of about 79 per cent by volume or 76 per cent by weight of the air. In chemical
combination with other elements it occurs in sodium nitrate (Chile saltpetre), in proteins,
and many other animal and vegetable products.
Nitrogen can be prepared from the atmosphere (air) and also from the ammonium
nitrate in the Laboratory. Nitrogen is obtained in industry by the fractional distillation of
liquid air. This principal use of nitrogen is in the manufacture of ammonia (industrial
scale by Haber process), from which nitrogenous fertilizers, nitric acid and urea are now
mainly prepared.
There are three common oxides of nitrogen. They are dinitrogen oxide or nitrous
oxide, N2O; nitrogen oxide or nitric oxide, NO; and nitrogen dioxide, NO 2 or dinitrogen
tetroxide, N2O4, depending on the conditions.
Nitric acid is manufactured by the Ammonia-Oxidation Process. Nitric acid is a
colourless, fuming liquid. Nitric acid is used mainly of the manufacture of explosives and
dyes. A valuable fertilizer can be obtained by neutralizing the acid with lime.

Questions and problems

1. How would you prepare nitrogen from
(a) the air (b) an ammonium salt?
What difference, if any, would there be, between the two samples so prepared?

2. Draw a fully labelled diagram and give an equation to show how you would prepare a
dry sample of ammonia in the laboratory starting from a named ammonium salt and a
named alkali.

3. Describe the Haber Process for the manufacture of ammonia.

4. Write equations for the following reactions. When ammonia is passed

(i) into dilute sulphuric acid (ii) over heated copper (11) oxide.

5. In the laboratory, a .steady stream of ammonia can be prepared by the reaction

between ammonium sulphate and an alkali.
(a) Name the alkali you would use.
 173

(b) State whether it is necessary to heat the mixture.

(c) Write the equation for the necessary to heat the mixture.
(d) Name the reagent you would use to dry the gas and explain your choice.
(e) State how you would collect a sample of dry ammonia.

6. Ammonia is manufactured by the Haber Process.

(a) Name the sources from which the starting materials are obtained.
(b) In what ratio by volume, should the nitrogen and hydrogen be mixed? Explain
why this ratio is used.
(c) What is the purpose of the catalyst?
(d) Is there complete conversion of all the nitrogen and hydrogen to ammonia?
(e) How can the ammonia be separated from the uncombined nitrogen and hydrogen?
(f) How are the uncombined nitrogen and hydrogen left after removal of the ammonia
treated?

7. Give the name, formula and colour of three gaseous oxides of nitrogen. Which of
these gases is very soluble in water? Outline briefly the preparation of this oxide?

8. Which of the oxides of nitrogen could be confused with oxygen ? Why ? By what test
can you distinguish between a jar of this gas and a jar of oxygen?

9. If you are provided with concentrated nitric acid, copper, and water but no other
chemical substances, describe briefly the reactions you would use prepare nitrogen
oxide and nitrogen dioxide.

10. Nitric acid can be prepared by the action of sulphuric acid on sodium nitrate.
(a) Give me equation for the reaction.
(b) State whether the sulphuric acid used must be concentrated or dilute, and whether
heat is needed.
(c) Give a diagram of the apparatus you would set up for the preparation.
(d) The nitric acid prepared is usually a yellow liquid. Why is it so?

11. (a) Describe the manufacturing process of nitric acid from ammonia.
(b) Describe two reactions each in which nitric acid acts
(i) as an acid (ii) as an oxidizing agent

12. Outline the chemical reactions involved in the synthesis of nitric acid starting from
nitrogen.

13. Write equations to show the formation of five different nitrogenous products in the
reduction of nitrate ions in acid solution.

14. Describe and explain the action of nitric acid of varying concentrations on
magnesium.
 174

15. Write equations to illustrate the action of heat on ammonium nitrate, calcium nitrate,
sodium nitrate and mercury (II) nitrate.

16. Give a chemical test which would enable you to decide if a given colourless solution
contained the nitrate ion.

17. A concentrated acid. A, is added to copper turnings and water in a flask and a brown
gas, B, is seen to fill the flask. In a second experiment the air is removed by filling the
flask with nitrogen before the acid is added. It is found that a colourless gas, C, is
formed. A fresh portion of acid. A, is diluted and neutralized with potassium
hydroxide solution. The solution is carefully evaporated to dryness to give a white
solid, D. On heating D a new solid, E, and a gas, F, are formed. F is found to ignite a
glowing splint.Identify A, B, C, D, E, F and give equations for the reactions.

18. Nitrogen is necessary for all plants growth.

(a) Name a plant which can assimilate (take up) nitrogen from the atmosphere.
(b) Give the chemical name of a nitrogen-containing compound which is used as a
fertilizer.
(c) Explain how the soil obtained nitrates by natural means.

19. Write TRUE or FALSE for each of the following statements.

(a) Nitrogen is present to the extent of about 79 percent by volume of the air.
(b) Nitrogen gas is relatively active and support combustion.
(c) The principal use of nitrogen is in the manufacture of ammonia.
(d) Nitrogen can be manufactured by the Haber process.
(e) Ammonia gas turns moist red litmus paper blue.

20. Fill in the blanks with a suitable words or units or phrases.

(a) Ammonia gas has a characteristic .................... smell.
(b) A mixture of calcium cyanamide and carbon is known as ..................
(c)The considerable solubility of ammonia is after demonstrated in the laboratory by
the ..................... experiment.
(d) Dinitrogen oxide is used as anesthetic for minor operation and is given the
name ................... gas.
(e) Nitric acid can be prepared by heating potassium nitrate crystals with concentrated
.................

21. Select the correct word, or words given in the brackets.

(a) Ammonia is prepared in the laboratory by heating ammonium chloride with
(sodium oxide, calcium hydroxide, slaked lime).
(b) Nitrogen is obtained in industry by (destrictive, fractional, vacuum) distillation of
liquid air.
(c) Nitrogen combines with many metals on heating to a dull red heat or higher
temperatures to form (oxides, nitrates, nitrides).
 175

(d) The principal use of nitrogen is in the manufacture of (ammonia, urea, nitric acid).
(e) When nitric oxide comes in contact with oxygen of the air (white, brown, red)
fumes are formed.

22. Match each of the items given in List A with the appropriate item in List B.
List A List B
(a) Nitric acid neutralizes bases (i) Haber process
(b) Dinitrogen oxide is also known as (ii) Forming metallic nitrates
(c) The manufacturing method of ammonia (iii) Nitrogen monoxide
(d) The solubility of ammonia can be tested (iv) Laughing gas
(e) Reddish brown fumes are formed when (v) Fountain experiment
a gas comes in contact with oxygen of
the air

23. What would be a suitable drying agent for ammonia gas ?

(A) Calcium chloride (B) Concentrated sulphuric acid
(C) Phosphorus (V) oxide (D) Calcium oxide

24. A gas was neutral, colorless and did not support combustion. Which gas might it
be ?
(A) Ammonia (B) Oxygen (C) Nitrogen dioxide (D) Nitrogen

25. Ammonia gas is collected by the ‘ upward delivery ’ when prepared in the
laboratory. The reason for this is because
(A) it is an extremely soluble gas (B) it is alkaline
(C) it is lighter than air (D) it is poisonous

26. Which of the following when heated would not give off ammonia ?
(A) Ammonium nitrate and sodium hydroxide
(B)Sodium nitrate and potassium hydroxide
(C) Ammonium sulphate and calcium hydroxide
(D) Sodium hydroxide and ammonium sulphate

27. When a gas X is mixed with hydrogen chloride gas, a dense white smoke forms.
Gas X is therefore : (A) carbon dioxide (B) ammonia (C) oxygen
(D) nitrogen

28. The process by which atmospheric nitrogen is turned into nitrate in the soil is called
(A) nitration (B) fixing (C) neutralizing (D) reduction
 176

29. When aqueous ammonia was added to a solution of a metal salt, a reddish-brown
precipitate formed. This indicates the presence of the metal ion :
(A) Ca 2+ (aq) (B) Cu + (aq) (C) Fe 2+ (aq) (D) Fe 3+ (aq)

30. You are given the following gases and answer the following questions given
below. nitrogen, ammonia, nitrogen oxide, dinitrogen oxide, nitrogen dioxide
(a) Which of the gases is neutral oxide ?
(b) Which of the gases is very soluble in water. Name the experiment to
confirm its solubility.
(c) Which of the gases is used in minor surgical operations ?
(d) Which of the gases solidify at – 10oC and under normal pressure ?
(e) Which of the above gases will give a brown ring when reacted with
iron(II) sulphate ? Give the formula of the brown ring compound.

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