Group 15 elements

General Introduction:

1) Group 15 of the periodic table includes five elements namely nitrogen (N).
phosphorus (p). Arsenic (As). antimony (Sb) and bismuth (Bi).
2) This group is regarded as nitrogen family or VA sub group elements of 15 group
elements.
3) There is a transition from non-metallic to metallic character as we go down the
group. Nitrogen and phosphorous are non-metals: arsenic and antimony are
semimetals or metalloids: bismuth is metallic.
4) Besides nitrogen the other important elements of group 15 is phosphorous they
form a group which exemplifies a gradual change in physical and chemical
properties.
5) Their physical properties like physical states, allotropic modifications, electronic
configuration, oxidation states etc vary gradually from N 2 to Bi.

Electronic configuration:-
The valence shell electronic configuration of these elements is ns2 np3. The s
orbital in these elements is completely filled and P-orbitals are half-filled, making
their configuration extra stable.

Electronic configuration of VA group elements

S.no Element Symbol Atomic Gondensed representation

number
of configuration
2 3
1 Nitrogen N 7 [He] 2 s 2 p
2
2 Phosphorous P 15 [Ne] 3s 3p3
3 Arsenic As 33 [Ar] 3d10 4s2 4p3
4 Antimony Sb 51 [Kr] 4d105s25p3
5 Bismuth Bi 83 [Xe] 4f14 5d106s26p3

Occurrence:
Occurrence of N:
Molecular nitrogen comprises 78% by volume of the atmosphere. In the earth’s
crust, it occurs as sodium nitrate. NaNO3 [called Chile saltpeter] and potassium
nitrate [Indian saltpeter]. It is found in the form of proteins in plants and animals.

Occurrence of P:
Phosphorus occurs in minerals of the apatite family. Ca9(PO4)6.CaX2 [X = F. Cl or
OH].

Example: Florapatite:
3Ca3(PO4)2 CaF2
Chlorapatite: 3Ca3(PO4)2 CaCl2
Phasphorite rock. Ca3(PO4)2

Occurrence of As, Sb & B:

They are not very abundant. Mostly they are available as their sulphide.
The extent of their occurrence in nature is given in the table given below

Element : N P As Sb Bi

Abundance in earths curst: 19 11.201.8 0.2 0.008 ppm

Oxidation states:
As per their configurations. these elements may utilize either the three electrons in the p-
orbitals or five electrons present in both s and p-orbitals in the exhibition of oxidation
states in their chemical compounds they exhibit common oxidation states +III and +V.
But nitrogen, being the first element of the group, is an exception in this respect. It
exhibits a number of oxidation states besides the two common oxidation states. The
oxidation states shown by nitrogen are shown with suitable examples, in the table given
below.
Oxidation states of nitrogen

S.No Name of the I- ormula of the Oxidation state

compound compound of Nitrogen
1 Ammonia NH3 -III
2 Hydrazine N2H4 - II
3 Hydroxylamine NH2OH -I
4 Nitrous oxide N2O +I
5 Nitric oxide NO + II
6 Nitrous acid HNO2 + III
7 Nitrogen dioxide NO + IV
8 Nitric acid HNO3 +V
9 Hydrazoic acid N3H - 1/3

Nitrogen, being more electronegative than hydrogen, exhibits negative oxidation states in
its binary compounds with hydrogen.

Phosphorus, similar to nitrogen, exhibits all possible oxidation states between

— III and + V in its hydrides, oxides or oxoacids. The oxidation states shown by
phosphoros are shown with suitable examples, in the table given below.

S.no Name of the compound I-ormula of the Oxidation state

compound of phosphrous
1 Galcium phosphide Ca3P2 - III
2 Phosphorous trichloride PCI3 + III
3 Phousphorous penta PCI5 +V
chloride
4 Hypophos phorous acid H3PO2 +1

Arsenic, Anitmony and Bismuth are either trivalent or pentavalent in their common
compounds. The stability of +V state decreases as one moves down the group from
nitrogen to Bismuth. The heavy element bismuth is stable in +III oxidation state. This is
due to inert pair effect of ns - electrons.
Physical characteristics of Group 15 elements:

The important physical constants of group 15 elements are given below.

Physical property N P As Sb Bi
Atomic radius(pm) 70 110 120 140 150
Ionic raoius(pm) 171(N3-) 212 (p3-) 222 (AS3-) 76(Sb3+) 103(Bi3-)

ionisazion energy IE1 1042 1012 947 834 703

(kj mol-) IE2 2856 1903 1798 1594 1610
IE3 4577 2910 2736 2443 2446
Electronagitivity 3.0 2.1 2.0 1.9 1.9
Melting point (k) 63 317.1 1089 904 544.4
Boiling point (k) 77 553.5 888 1860 1837
Density (g cm-3) 0.879 1.823 5.778 6.697 9.808

The important physical characteristics are discussed below:

Atomic and ionic radii:

On going down the group, the atomic radius increases due to the increase in number of
shells.

Melting and boiling points:

Melting points (except for antimony and bismuth) and boiling points increase on going
down the group from N to Bi.

Ionization energies:
The first ionization energy of the group 15 elements are larger due to stable electronic
configuration.
The valence shell electronic configuration of atoms of group 15 is ns2. npx1 npy1, npz1 and
are stable.
Therefore, they have high ionization energies. On going down the group, the ionization
energies decrease, this is due to increase in atomic size.

Electro negativity:
The elements of group 15 have smaller size and greater nuclear charge of atoms and
therefore they have higher electro negativity value. This is due to increase in size of the
atoms and shielding effect of inner electron shells on going down the group.

Metallic character:
The elements of group 15 are less metallic. However on going down the group, the
metallic character increases from N to Bi

Example:
a. N and P are non- metals
b. As and Sb are partly non-metals or metalloids
c. Bi is a metal.

Catenation:
The elements of group 15 also show a tendency to form bonds having the atoms of the
same elements known as catenation. All these elements show this property but to a much
 smaller extent than carbon. This tendency is maximum in nitrogen in this group elements.
This is least in Bi.

Example:
Hydrazine [H2N – NH2] two nitrogen atom chain.
Hydrazoic acid [N3H] three nitrogen atom chain.
Diphosphine [P2H4] two phosphorus atom chain.
On going down the group, the catenation capacity among the group 15 elements
gradually decreases.

Allotropy:

We know that from the discussion in earlier chapters on B and C the property of an
element to exist in two or more physical forms having more or less similar chemical
properties but different physical properties" is known as allotropy

Except Bismuth, all other elements of this group show allotropy.

For example:

I. Solid N2 exists in -alpha and -beta nitrogen forms.

II. Phosphorous exists in many allotropic forms such as white P;

RedP; scarlet P; a -Black P; 8-black P and violet P.

III. Arsenic exist as — yellow or grey arsenic

IV. Antimony exist as — yellow or silver grey

V. Only one form of bismuth is known.

Chemical properties:
a. Reactivity towards hydrogen:
All the elements of group 15 form hydrides of the type EH3 where E = N. P. As. Sb or Bi.
The hydrides are.
1. Ammonia NH3

2. Phosphine PH3
AsH3
3. Arsine
SbH3
4. Stibine
BiH3
5. Bismuthine
They are all volatile. Covalent compounds. The thermal stability decreases down the group from
Ammonia to Bismathine. Their reducing nature increases gradually.

b. Reactivity towards oxygen:

All the elements form two types of oxides. E 2 O 3 and E 2 O 5 The oxides of nitrogen (N 2 O 3) and
phosphorus (P 2 O 3) are purely acidic. Those of arsenic and antimony are amphoteric and those that of
bismuth predominantly basic.
c. Reactivity towards halogens:
These elements react to form two series of halides. Ex3 and Ex5 Nitrogen does not form pentahalide
due to non-availability of the d orbitals in its valence shell. Pentahalides are more covalent than
trihalides.
d. Reactivity towards metals:
All these elements react with strongly positive metals to form binary compounds exhibiting -3
oxidation state..

Example: Ca3N 2 [Calcium nitride], Ca 3P 2 [Calcium phosphide], Na3As [Sodium arsenide], Znt3Sb 2
[Zinc antimonide] and Mgt 3B i2 [Magnesium bismuthide]

Nitrogen is the most abundant of all the elements of groups 15.

1. Preparation of nitrogen:
a. Laboratory preparation of nitrogen:
In the laboratory, nitrogen is prepared byheating ammonium nitrite. Ammonium nitrite is
decomposes to produce nitrogen gas. However, this reaction is very fast and may prove to be explosive.
Therefore in the lab a mixture of ammonium chloride and sodium nitrite is heated with small amount
of water added to it.

The mixture contains ammonium chloride and sodium nitrite approximately in the ratio of 4:5 by mass.
The presence of water prevents ammonium chloride from subliming when heated. Initially two
substances undergo double decomposition to form sodium chloride and ammonium nitrite

The ammonium nitrite so formed then decomposes to form nitrogen gas and water vapor.
 NH4NO2 2H2O N2
aq
g

Nitrogen is collected by downward displacement of water because N2 is insoluble in water.

Laboratory preparation of nitrogen

The nitrogen obtained from nitrogenous compounds like NH4NO2is called “Chemical nitrogen". As it is
free from other atmospheric gases, it has a lower vapour density compare to the nitrogen obtain from the
atmosphere

b. Preparation of nitrogen from chemical compounds:

(i) By passing ammonia over heated metallic oxides like copper oxide and lead oxide. NH3 is oxidized
to N2

(ii) NH3can also be oxidized by bleaching powder to N2

(iii) Ammonium dichromate. when heated decomposes with evolution of heat. The products of
decomposition are a green coloured chromic oxide. solid water vapor and nitrogen gas.
 II. Properties:

Physical properties:
It is a colourless. tasteless. odourless gas. slightly soluble in water. It is lighter than air. (V.D of
nitrogen = 14. V.D of air = 14.4)

Chemical properties:
Nitrogen does not easily combine with other elements under ordinary conditions
The triple bond between the N - atoms is so strong. that a large amount of energy is needed to break
this bond. Hence special conditions such as temperature. pressure. catalyst. promoters. high voltage
etc. are needed to make nitrogen combine with other elements.

i. With hydrogen:
Nitrogen combines with hydrogen in presence of electric sparks to form ammonia. Usually ammonia
is obtained by treating a 3:1 mixture (by volume) of hydrogen and nitrogen at about450’c. and a
pressure of 200 to 1000 atm. in the presence of finely divided iron as catalyst and molybdenum as
promoter.

ii. With oxygen:

Nitrogen combines with oxygen only under the influence of an electric arc at a temperature of
3000’c - 5000’c to form nitric oxide. Also when lighting takes place in the sky form nitric oxide.

iii. With metals:

Red hot or burning magnesium combine with nitrogen to form magnesium nitride.
 Burning of magnesium in nitrogen

iv. With calcium carbide:

When N2 gas is passed over hot calcium carbonate (at 800 oC to 1000 oC). Calcium cynamide an
important fertilizer is formed. The commercial name of calcium cynamide is nitrolim.

CaC2 N2 CaCN2 C
800-1000 CO
Nitrolium

Summary of chemical properties of nitrogen

Uses of Nitrogen:
1. Nitrogen is used in high temperature thermometers where mercury cannot be used. A
volume of nitrogen is enclosed in a vessel and introduced into the region of high
temperature depending upon the temperature. expansions in the nitrogen gas 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 giving a longer life to the filament of a bulb.
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 micro organisms
do not grow.
4. It is also used in metal working operating 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 in petroleum and paint industries to provide inactive atmosphere to prevent
fires or explosions.
7. It is used in the industrial preparation of ammonia. by Haber’s process. NH3 is converted
into ammonium salts, nitric acid, urea, calcium cynamide fertilizers etc.
8. Liquid nitrogen is used as refrigerant for food, for storage of blood, cornea etc.
9. Liquid nitrogen is used in scientific research especially in the field of super conductors.
10. Nitrogen is essential for synthesis of proteins in plants.
11. Liquid nitrogen is used in oil fields. to extinguish oil fires.
Ammonia:
Ammonia is by far the most important hydride of the V group elements. Ammonia is
present in atmospheric air and is present in the trace amounts. However in sewage water, it is present
in greater quantities. It is released into the atmosphere by the decomposition of organic wasters of
humans.
(A) Preparation:
a. Laboratory preparation of ammonia:
i. In the laboratory, ammonia is usually prepared by heating a mixture of ammonium chloride and
slaked lime taken in the ratio of 2:3 by mass. The arrangement of the apparatus is shown in the below
figure.

Preparation of ammonia in the laboratory from ammonium chloride

2 NH4Cl + Ca(OH)2 CaCl2 + 2H2O + 2NH3
i. As ammonia is lighter than air, it is collected by the downward displacement of air.
ii. Take magnesium nitride in a conical flask as shown in the below figure.
Add hot water to it through the thistle funnel. The respective hydroxide and ammonia are formed.

Mg3N2 + 6H2O 3 Mg(OH)2 + 2NH3

b. Industrial preparation of Ammonia:

Haber’s process:
On large scale, ammonia is manufactured by haber’s process starting with nitrogen and hydrogen.
Reactants: Nitrogen gas – 1 volume and hydrogen gas – 3 volumes.
Reaction:
N2(g) + H2(g) 2NH3(g) + heat
Conditions:
Temperature Pressure Catalyst Promoter
4500C-5000C 200-300 atm Iron Molybdeum

Manufacture of Ammonia – Haber’s Process

The reaction in Haber’s process is exothermic and so external heating not required once the reaction
starts. Lowering the temperature to 4500C-5000C favors the reaction, but lowering the temperature
below 4500C- 5000C brings the yield.
Properties:
a. Physical properties:
i. Colour: Ammonia is a colourless gas.
ii. Odour: It has a characteristic pungent odour.
iii. Taste: It is bitter to taste.
iv. Solubility: Ammonia is one of the most soluble gasses in water. At 00C and 760 mm of Hg
pressure on e volume of water can dissolve nearly 1200 volumes of ammonia.
b. Chemical properties:
1. Reaction with oxygen:
i) Burning of ammonia in oxygen:
When ammonia is burnt in oxygen or in air, it catches fire and the following reaction takes
place.
4NH3 + 3O2 6H2O + 2N2
 Burning of ammonia in oxygen
ii) Catalytic oxidation of Ammonia:
The platinum coil is heated at 800 oC in a combustion tube till it becomes white. Then ammonia and
oxygen are passed through the tube. Under these conditions and in the presence of P catalyst ammonia
combines with free oxygen to form nitric oxide and water vapour. This reaction is used in the
manufacturing of HNO3 by Ostwald’s process.
4NH3 + 5O2 4NO + 6H2O

2. Reaction with water:

Ammonia gas is highly soluble in water and form ammonium hydroxide
NH3 + H2O NH4OH
Highly concentrated ammonia solution is liquor ammonia.
3. Reaction with acids:
Ammonia reacts with the acids to form their respective ammonium salts
2NH3 + H2SO4 (NH4)2 SO4
Ammonium sulphate similarly HNO3 gives NH4NO3 and HCl gives NH4Cl.
4. Reaction with chlorine:
Chloride reacts with NH3 in which the products formed vary with the experimental conditions.
i. With excess of chlorine, nitrogen trichloride and HCl are formed.
NH3 + 3Cl2 NCl3 + N2
ii. With excess of ammonia, nitrogen and ammonium chloride are formed.
8NH3 + 3Cl2 6NH4Cl + N2
5. Reaction with aqueous solutions of salts:
When ammonia is passed through aqueous solutions of certain salts, the respective metallic
hydroxides are precipitated.
ZnSO4 + 2NH3 + 2H2O (NH4)SO4 + Zn(OH)2
(white gelatinous ppt)
FeCl3 + 3NH3 + 3H2O 3 NH4Cl + Fe(OH)3
(Reddish brown ppt)
6. Reaction with heated metallic oxide:
When passed over heated metallic oxides are reduced to the corresponding metal.
Example: 3CuO + 2NH3 3Cu + 3H2O + N2
The chemical properties of ammonia are summarized in the given figure shown below.

Nitric acid:
Nitrogen forms oxo acids such as H2N2O (hypo nitrous acid),
HNO2 (nitrous acid) and HNO3 (nitric acid) among them HNO3 is the most important.
Nitric acid is present in small quantities in the atmosphere. But during thundershowers, a large
quantities of this acid is produced in the atmosphere. It then comes down as a very dilute solution in
rain water. The nitric acid present in the rain reacts with the minerals present in the soil and gets
converted into nitrates.
A. Preparation:
a. Laboratory preparation of Nitric acid:
In the laboratory nitric acid is prepared by distilling a mixture of concentrated sulphuric
acid and potassium nitrate or sodium nitrate. The apparatus is set up as shown in figure given below.

Preparation of nitric acid in the laboratory

The reaction is
NaNO3 + H2SO4 NaHSO4 + HNO3
Nitric acid being strong corrosive, attacks rubber cork etc.
b. Manufacture of nitric acid by Ostwald’s process:
On a large scale it is prepared mainly by Ostwald’s process. This method is based upon catalytic
oxidation of NH3 by atmospheric oxygen.
Reactants: Pure dry ammonia (1 volume) and air (10 volumes)
Reactions:
Catalytic oxidation of ammonia to form nitric oxide.
4NH3 + 5O2 4NO + 6H2O + 21.5 K.cal
Oxidation of nitric oxide to nitrogen dioxide.
2 NO + O2 2NO2
Absorption of nitrogen dioxide in water to give nitric acid.
4NO2 + 2H2O + O2 4HNO3

Manufacture of HNO3 – Ostwald’s process

B. Properties:
a. Physical properties of nitric acid:
i. Colour: Pure nitric acid is colorless but commercial nitric acid may be yellow & brown, due to
presence of dissolved nitrogen dioxide.
ii. Odour: Nitric acid is fuming, hygroscopic liquid, the fumes of which give it a choking smell.
b. Chemical properties of nitric acid:
1. Stability:
Pure nitric acid is not very stable. Even at ordinary temperature, in presence of sunlight it
undergoes slight decomposition. That is why it is yellow in colour. As the temperature increases, the
rate of decomposition also increases. On strong heating it decomposes completely to give nitrogen
oxide, water and oxygen.
4HNO3 2H2O + 4NO2 + O2
2. Reaction with Gold and Platinum:
Gold and platinum do not dissolve in either dil or conc. nitric acid. But they dissolve in “Aqua
Regia” it is a mixture of concentrated nitric acid and concentrated hydrochloric acid in the proportion
of 1:3 by volume.
3HCl + HNO3 2H2O + NOCl + Cl2
3. With Non-metals:
With hot concentrated nitric acid, non-metals are oxidized to their oxides while the acid itself is
reduced to nitrogen dioxide.
6HNO3 + S 2H2O + H2SO4 + 6NO2

Summary of properties of nitric acid

Oxides of Nitrogen and Allotropes of Phosphorous

A. Oxides of Nitrogen:
Nitrogen forms a wide range of oxides.Of these oxides, N2O and NO are neutral oxides. The higher
oxides N2O3, N2O4 and N2O5 are acidic. Lewis dot main resonance structures oxides are given in table
given below.
Structures of oxides of Nitrogen
B. Allotropes of Phosphorus:
Phosphorus is found in many allotropic forms, the important ones being white, red and black.
a. White phosphorus or Yellow phosphorus:
Elemental phosphorus is obtained by heating phosphate rock will coke and silica in an electric furnace
at about 1770 K. The phosphorus so formed is White phosphorus. It is collected under water.
The reactions may be represented as
2Ca(PO4)2 + 6SiO2 6CaSiO3 + P4O10
P4O10 + 10C 4P + 10 CO
i. It is a soft waxy solid with garlic smell.
ii. It is a poisonous in nature.
iii. It turns yellowish on exposure to light. For this reason. it is also called
yellow P.
iv. It is not soluble in water but soluble in carbon disulphide.
v. It undergoes spontaneous combustion in air and produces greenish
glow.
vi. It exists as P4 molecules both in solid and vapor state. The four atoms
in P4 molecule occupythe corners of regulartetrahedron. P4 may be
polymeric.

b. Red Phosphorus:
It is prepared by heating white phosphorus to about 540 K in an inert atmosphere of nitrogen for
several hours.

White phosphorous Red phosphorous

i. It is a hard crystalline solid without any smell.

ii. It is non- poisonous in nature
iii. It has a layer type structure in which each layer consists of
phosphorus atoms.
iv. It is more stable and relatively less reactive.
v. It consists of tetrahedral units of P4 linked to one anotherto
constitute linear chains.

C. Black Phosphorus:
It is prepared by heating white phosphorus to about 470 K under high pressure of 1200 atmospheres in
an inert atmosphere.
White phosphorous Black phosphorous
i. It has metallic lustre.
ii. It is most inactive form of phosphorus.
iii. It has a layer type structure in which each layer consists of
phosphorus atoms.

Some physical properties of three forms of Phosphorus are given below.

Property White phosphorus Red phosphorus Black phosphorus

colour White, but turns yellow on Dark red Black
exposure
State Waxy solid, can be cutwith Brittle powder Crystalline with greasy
knife touch
Smell Garlic smell Odorless ---
Density 1.84 2.1 2.69
Ignition 307 K 533 K 673 K
temperature
Melting point 317 K Does not melt 860 K

Elementary ideas of oxoacids, super phosphate of lime

A. Oxo acids of Nitrogen:

The elements of group 15 form a large number oxoacids.

i. Oxo acids of nitrogen:

Nitrogen forms a number of oxoacids. The important oxoacids of nitrogen with their formulae. one method of
preparation are given in table.

Formula Preparation
Name Oxidation state
of Nitrogen

Hyponitrous H2N2O2 or +1 (Silver hyponitrite ) +

acid HNO anhydrous HCl
 Nitrous acid HNO2 +3 (Barium nitrite) + ice cold
H2SO4

(Potassium nitrate)+ conc.

Nitric acid HNO3 +5
H2SO4

Out of this oxoacids of nitrogen, nitric acid is the most important. It is very strong oxidising agent and is quite useful.
B. Oxoacids of phosphorus:
Phosphorus forms a number of oxoacids. The important oxoacids of phosphorus with their formulas,
one methods of preparation and the presence of some characteristic bonds in their structures are given
in below table.

Name Formula Oxidation state Preparation

of Phosphorus
Hypo phosphorous acid H3PO2 +1 White P + alkali
Ortho phosphorous acid H3PO3 +3 P 2 O3 + H 2 O
Pyrophosphorus acid H4 P 2 O5 +3 PCI3 + H3PO3
Ortho phosphoric acid H3PO4 +5 P4O10+ H2O
Pyro phosphoric acid H4 P 2 O7 +5 Heat phosphoric acid
Hypophosphoric acid H4 P 2 O6 +4 Red P + alkali
Meta phosphoric acid HPO3 +5 H3PO3 + Br2 heated in a
sealed tube.

Among the oxoacids of phosphorus. ortho phosphoric acid is the most important and used in the
manufacture of phosphate fertilizers.

The structures of oxoacids of phosphorus and their structural similarities are shown below.
The formulas of oxoacids of P can be remembered as:
a. Meta acid is used for the acid obtained by the loss of one water
molecule.
b. Pyro acid is used forthe acid obtained by heating two molecules with
loss of one water molecule.
c. Hypo is generally forthe acid having lower oxygen contentthan the
parent acid.
However, metaphosphoric acid does not exist as simple monomer; it exists as cyclo meta phosphoric
acid or poly meta phosphoric acid.
C. Superphosphate of lime:
i. Preparation of superphosphate of lime:
Superphosphate of lime is also known as calcium superphosphate. This is one ofthe few water soluble phosphates. A
mixture of calcium dihydrogen phosphate [Ca (H2 PO4)2] and gypsum (CaSO4.2H2) is known as superphosphate of lime.

Superphosphate is made by treating well powered phosphate rock (calcium phosphate) with calculated quantities of
commercial concentrated sulphuric acid (chamber acid).

Ca3 (PO4)2 + 2H2 SO4 + 4H2O Ca (H2PO4)2 + 2 [CaSO4. 2H2O]+ Heat

Industrial method:
The plant used for the preparation ofsuperphosphate of lime is shown in figure given below

Manufacture of Superphosphate of lime

1 . St ir r er
2 . Cha mber ac id
3 . Valves
4 . Brick work Dens
5 . Waste gases

The phosphate rock is ground to a fine powder. It is charged into a cast iron mixer. A calculated quantity of Conc. H2SO4
is added and the reaction mixture is stirred with blades present in the mixer. When the reaction has started it is dumped into one
of the two dens. D1 and D2through either the value V1 or through the value V2. The reaction is allowed to take place for
24-36 hours in the dens. In this period temperature rises to about 373-383 K. The carbonate and the fluoride impurities in the
phosphate rock react with H2SO4 and liberate CO2 and HF gases. They escape through the outlet at the top. The final
product is a hard mass and is sold as superphosphate of lime.