170Chemistry
In Class XI, you have learnt that the p-block elements
are placed in groups 13 to 18 of the periodic table.
Their valence shell electronic configuration is ns
2
np
1–6
(except He which has 1s
2
configuration). The properties
of p-block elements like that of others are greatly
influenced by atomic sizes, ionisation enthalpy, electron
gain enthalpy and electronegativity. The absence of d-
orbitals in second period and presence of d or d and f
orbitals in heavier elements (starting from third period
onwards) have significant effects on the properties of
elements. In addition, the presence of all the three types
of elements; metals, metalloids and non-metals bring
diversification in chemistry of these elements.
Having learnt the chemistry of elements of Groups
13 and 14 of the p-block of periodic table in Class XI,
you will learn the chemistry of the elements of
subsequent groups in this Unit.
The
p
-Block
Elements
7
The
p
-Block
Elements
After studying this Unit, you will be
able to
appreciate general trends in the
chemistry of elements of groups
15,16,17 and 18;
learn the preparation, properties
and uses of dinitrogen and
phosphorus and some of their
important compounds;
describe the preparation,
properties and uses of dioxygen
and ozone and chemistry of some
simple oxides;
know allotropic forms of sulphur,
chemistry of its important
compounds and the structures of
its oxoacids;
describe the preparation,
properties and uses of chlorine
and hydrochloric acid;
know the chemistry of
interhalogens and structures of
oxoacids of halogens;
enumerate the uses of noble
gases;
appreciate the importance of
these elements and their
compounds in our day to day life.
Objectives
Diversity in chemistry is the hallmark of p–block elements manifested
in their ability to react with the elements of s–, d– and f–blocks as
well as with their own.
Group 15 includes nitrogen, phosphorus, arsenic, antimony, bismuth
and moscovium. As we go down the group, there is a shift from non-
metallic to metallic through metalloidic character. Nitrogen and
phosphorus are non-metals, arsenic and antimony metalloids, bismuth
and moscovium are typical metals.
Molecular nitrogen comprises 78% by volume of the atmosphere.
In the earth’s crust, it occurs as sodium nitrate, NaNO
3
(called Chile
saltpetre) and potassium nitrate (Indian saltpetre). It is found in the
form of proteins in plants and animals. Phosphorus occurs in minerals
7.17.1
7.17.1
7.1
Group 15Group 15
Group 15Group 15
Group 15
ElementsElements
ElementsElements
Elements
7.1.1 Occurrence
UnitUnit
UnitUnit
Unit
7
2020-21
171 The p-Block Elements
of the apatite family, Ca
9
(PO
4
)
6
. CaX
2
(X = F, Cl or OH) (e.g., fluorapatite
Ca
9
(PO
4
)
6
. CaF
2
) which are the main components of phosphate rocks.
Phosphorus is an essential constituent of animal and plant matter. It
is present in bones as well as in living cells. Phosphoproteins are present
in milk and eggs. Arsenic, antimony and bismuth are found mainly as
sulphide minerals. Moscovium is a synthetic radioactive element. Its
symbol is Mc, atomic number 115, atomic mass 289 and electronic
configuration [Rn] 5f
14
6d
10
7s
2
7p
3
. Due to very short half life and
availability in very little amount, its chemistry is yet to be established.
Here, except for moscovium, important atomic and physical
properties of other elements of this group along with their electronic
configurations are given in Table 7.1.
Table 7.1: Atomic and Physical Properties of Group 15 Elements
a
E
III
single bond (E = element);
b
E
3–
;
c
E
3+
;
d
White phosphorus;
e
Grey
α
-form at 38.6 atm;
f
Sublimation temperature;
g
At 63 K;
h
Grey
α
-form; * Molecular N
2
.
Trends of some of the atomic, physical and chemical properties of the
group are discussed below.
The valence shell electronic configuration of these elements is ns
2
np
3
.
The s orbital in these elements is completely filled and p orbitals are
half-filled, making their electronic configuration extra stable.
Covalent and ionic (in a particular state) radii increase in size
down the group. There is a considerable increase in covalent radius
from N to P. However, from As to Bi only a small increase in
covalent radius is observed. This is due to the presence of
completely filled d and/or f orbitals in heavier members.
Ionisation enthalpy decreases down the group due to gradual increase
in atomic size. Because of the extra stable half-filled p orbitals electronic
configuration and smaller size, the ionisation enthalpy of the group 15
elements is much greater than that of group 14 elements in the
corresponding periods. The order of successive ionisation enthalpies,
as expected is
i
H
1
<
i
H
2
<
i
H
3
(Table 7.1).
7.1.2 Electronic
Configuration
7.1.3 Atomic and
Ionic Radii
7.1.4 Ionisation
Enthalpy
Property N P As Sb Bi
Atomic number 7 15 33 51 83
Atomic mass/g mol
–1
14.01 30.97 74.92 121.75 208.98
Electronic configuration [He]2s
2
2p
3
[Ne]3s
2
3p
3
[Ar]3d
10
4s
2
4p
3
[Kr]4d
10
5s
2
5p
3
[Xe]4f
14
5d
10
6s
2
6p
3
Ionisation enthalpy I 1402 1012 947 834 703
(
i
H/(kJ mol
–1
) II 2856 1903 1798 1595 1610
III 4577 2910 2736
2443 2466
Electronegativity 3.0 2.1 2.0 1.9 1.9
Covalent radius/pm
a
70 110 121 141 148
Ionic radius/pm 171
b
212
b
222
b
76
c
103
c
Melting point/K 63* 317
d
1089
e
904 544
Boiling point/K 77.2* 554
d
888
f
1860 1837
Density/[g cm
–3
(298 K)] 0.879
g
1.823 5.778
h
6.697 9.808
2020-21
172Chemistry
The electronegativity value, in general, decreases down the group with
increasing atomic size. However, amongst the heavier elements, the
difference is not that much pronounced.
All the elements of this group are polyatomic. Dinitrogen is a diatomic gas
while all others are solids. Metallic character increases down the group.
Nitrogen and phosphorus are non-metals, arsenic and antimony metalloids
and bismuth is a metal. This is due to decrease in ionisation enthalpy and
increase in atomic size. The boiling points, in general, increase from top to
bottom in the group but the melting point increases upto arsenic and then
decreases upto bismuth. Except nitrogen, all the elements show allotropy.
Oxidation states and trends in chemical reactivity
The common oxidation states of these elements are –3, +3 and +5.
The tendency to exhibit –3 oxidation state decreases down the group due
to increase in size and metallic character. In fact last member of the group,
bismuth hardly forms any compound in –3 oxidation state. The stability
of +5 oxidation state decreases down the group. The only well characterised
Bi (V) compound is BiF
5
. The stability of +5 oxidation state decreases and
that of +3 state increases (due to inert pair effect) down the group. Besides
+5 oxidation state, nitrogen exhibits + 1, + 2, + 4 oxidation states also
when it reacts with oxygen. However, it does not form compounds in
+5 oxidation state with halogens as nitrogen does not have d-orbitals
to accommodate electrons from other elements to form bonds.
Phosphorus also shows +1 and +4 oxidation states in some oxoacids.
In the case of nitrogen, all oxidation states from +1 to +4 tend to
disproportionate in acid solution. For example,
3HNO
2
HNO
3
+ H
2
O + 2NO
Similarly, in case of phosphorus nearly all intermediate oxidation
states disproportionate into +5 and –3 both in alkali and acid. However
+3 oxidation state in case of arsenic, antimony and bismuth becomes
increasingly stable with respect to disproportionation.
Nitrogen is restricted to a maximum covalency of 4 since only four
(one s and three p) orbitals are available for bonding. The heavier elements
have vacant d orbitals in the outermost shell which can be used for
bonding (covalency) and hence, expand their covalence as in PF
6
.
Anomalous properties of nitrogen
Nitrogen differs from the rest of the members of this group due to
its small size, high electronegativity, high ionisation enthalpy and
non-availability of d orbitals. Nitrogen has unique ability to form
p
ππ
ππ
π
-p
ππ
ππ
π
multiple bonds with itself and with other elements having
small size and high electronegativity (e.g., C, O). Heavier elements of
this group do not form p
π
-p
π
bonds as their atomic orbitals are so
large and diffuse that they cannot have effective overlapping.
Thus, nitrogen exists as a diatomic molecule with a triple bond (one
s and two p) between the two atoms. Consequently, its bond enthalpy
(941.4 kJ mol
–1
) is very high. On the contrary, phosphorus, arsenic
and antimony form single bonds as P–P, As–As and Sb–Sb while
bismuth forms metallic bonds in elemental state. However, the single
7.1.5
Electronegativity
7.1.6 Physical
Properties
7.1.7 Chemical
Properties
2020-21
173 The p-Block Elements
N–N bond is weaker than the single P–P bond because of high
interelectronic repulsion of the non-bonding electrons, owing to the
small bond length. As a result the catenation tendency is weaker in
nitrogen. Another factor which affects the chemistry of nitrogen is
the absence of d orbitals in its valence shell. Besides restricting its
covalency to four, nitrogen cannot form d
ππ
ππ
π
p
ππ
ππ
π
bond as the heavier
elements can e.g., R
3
P = O or R
3
P = CH
2
(R = alkyl group). Phosphorus
and arsenic can form d
ππ
ππ
π
d
ππ
ππ
π
bond also with transition metals when
their compounds like P(C
2
H
5
)
3
and As(C
6
H
5
)
3
act as ligands.
(i) Reactivity towards hydrogen: All the elements of Group 15
form hydrides of the type EH
3
where E = N, P, As, Sb or Bi.
Some of the properties of these hydrides are shown in Table
7.2. The hydrides show regular gradation in their properties.
The stability of hydrides decreases from NH
3
to BiH
3
which can
be observed from their bond dissociation enthalpy.
Consequently, the reducing character of the hydrides increases.
Ammonia is only a mild reducing agent while BiH
3
is the
strongest reducing agent amongst all the hydrides. Basicity also
decreases in the order NH
3
> PH
3
> AsH
3
> SbH
3
> BiH
3
. Due to
high electronegativity and small size of nitrogen, NH
3
exhibits
hydrogen bonding in solid as well as liquid state. Because of this,
it has higher melting and boiling points than that of PH
3
.
(ii) Reactivity towards oxygen: All these elements form two types
of oxides: E
2
O
3
and E
2
O
5
. The oxide in the higher oxidation state
of the element is more acidic than that of lower oxidation state.
Their acidic character decreases down the group. The oxides of
the type E
2
O
3
of nitrogen and phosphorus are purely acidic,
that of arsenic and antimony amphoteric and those of bismuth
predominantly basic.
(iii) Reactivity towards halogens: These elements react to form two
series of halides: EX
3
and EX
5
. Nitrogen does not form pentahalide
due to non-availability of the d orbitals in its valence shell.
Pentahalides are more covalent than trihalides. This is due to the
fact that in pentahalides +5 oxidation state exists while in the case
of trihalides +3 oxidation state exists. Since elements in +5 oxidation
Property NH
3
PH
3
AsH
3
SbH
3
BiH
3
Melting point/K 195.2 139.5 156.7 185
Boiling point/K 238.5 185.5 210.6 254.6 290
(E–H) Distance/pm 101.7 141.9 151.9 170.7
HEH angle (°) 107.8 93.6 91.8 91.3
f
H
V
/kJ mol
–1
–46.1 13.4
66.4 145.1 278
diss
H
V
(E–H)/kJ mol
–1
389 322 297 255
Table 7.2: Properties of Hydrides of Group 15 Elements
2020-21
174Chemistry
Though nitrogen exhibits +5 oxidation state, it does not form
pentahalide. Give reason.
Nitrogen with n = 2, has s and p orbitals only. It does not have d
orbitals to expand its covalence beyond four. That is why it does not
form pentahalide.
PH
3
has lower boiling point than NH
3
. Why?
Unlike NH
3
, PH
3
molecules are not associated through hydrogen bonding
in liquid state. That is why the boiling point of PH
3
is lower than NH
3
.
Preparation
Dinitrogen is produced commercially by the liquefaction and fractional
distillation of air. Liquid dinitrogen (b.p. 77.2 K) distils out first leaving
behind liquid oxygen (b.p. 90 K).
In the laboratory, dinitrogen is prepared by treating an aqueous
solution of ammonium chloride with sodium nitrite.
NH
4
CI(aq) + NaNO
2
(aq)
N
2
(g) + 2H
2
O(l) + NaCl (aq)
Small amounts of NO and HNO
3
are also formed in this reaction;
these impurities can be removed by passing the gas through aqueous
sulphuric acid containing potassium dichromate. It can also be obtained
by the thermal decomposition of ammonium dichromate.
(NH
4
)
2
Cr
2
O
7
Heat

N
2
+ 4H
2
O + Cr
2
O
3
Very pure nitrogen can be obtained by the thermal decomposition
of sodium or barium azide.
Ba(N
3
)
2
Ba + 3N
2
7.27.2
7.27.2
7.2
DinitrogenDinitrogen
DinitrogenDinitrogen
Dinitrogen
Example 7.1Example 7.1
Example 7.1
Example 7.1
Example 7.1
SolutionSolution
SolutionSolution
Solution
Example 7.2Example 7.2
Example 7.2Example 7.2
Example 7.2
SolutionSolution
SolutionSolution
Solution
state will have more polarising power than in +3 oxidation state,
the covalent character of bonds is more in pentahalides. All the
trihalides of these elements except those of nitrogen are stable.
In case of nitrogen, only NF
3
is known to be stable. Trihalides
except BiF
3
are predominantly covalent in nature.
(iv) Reactivity towards metals: All these elements react with metals
to form their binary compounds exhibiting –3 oxidation state,
such as, Ca
3
N
2
(calcium nitride) Ca
3
P
2
(calcium phosphide),
Na
3
As (sodium arsenide), Zn
3
Sb
2
(zinc antimonide) and Mg
3
Bi
2
(magnesium bismuthide).
Intext QuestionsIntext Questions
Intext QuestionsIntext Questions
Intext Questions
7.1 Why are pentahalides of P, As, Sb and Bi more covalent than their
trihalides?
7.2 Why is BiH
3
the strongest reducing agent amongst all the hydrides of
Group 15 elements ?
2020-21
175 The p-Block Elements
Preparation
Ammonia is present in small quantities in air and soil where it is
formed by the decay of nitrogenous organic matter e.g., urea.
On a small scale ammonia is obtained from ammonium salts which
decompose when treated with caustic soda or calcium hydroxide.
2NH
4
Cl + Ca(OH)
2
2NH
3
+ 2H
2
O + CaCl
2
(NH
4
)
2
SO
4
+ 2NaOH 2NH
3
+ 2H
2
O + Na
2
SO
4
Properties
Dinitrogen is a colourless, odourless, tasteless and non-toxic gas.
Nitrogen atom has two stable isotopes:
14
N and
15
N. It has a very low
solubility in water (23.2 cm
3
per litre of water at 273 K and 1 bar
pressure) and low freezing and boiling points (Table 7.1).
Dinitrogen is rather inert at room temperature because of the high
bond enthalpy of NN bond. Reactivity, however, increases rapidly with
rise in temperature. At higher temperatures, it directly combines with
some metals to form predominantly ionic nitrides and with non-metals,
covalent nitrides. A few typical reactions are:
6Li + N
2
Heat

2Li
3
N
3Mg + N
2
Heat

Mg
3
N
2
It combines with hydrogen at about 773 K in the presence of a
catalyst (Haber’s Process) to form ammonia:
N
2
(g) + 3H
2
(g)
773 k

2NH
3
(g);
f
H
y
= –46.1 kJmol
–1
Dinitrogen combines with dioxygen only at very high temperature
(at about 2000 K) to form nitric oxide, NO.
N
2
+ O
2
(g)
Heat

2NO(g)
UsesUses
UsesUses
Uses: The main use of dinitrogen is in the manufacture of ammonia and other
industrial chemicals containing nitrogen, (e.g., calcium cyanamide). It also
finds use where an inert atmosphere is required (e.g., in iron and steel industry,
inert diluent for reactive chemicals). Liquid dinitrogen is used as a refrigerant
to preserve biological materials, food items and in cryosurgery.
Write the reaction of thermal decomposition of sodium azide.
Thermal decomposition of sodium azide gives dinitrogen gas.
3 2
2NaN 2Na 3N
+
Example 7.3Example 7.3
Example 7.3Example 7.3
Example 7.3
SolutionSolution
SolutionSolution
Solution
Intext QuestionIntext Question
Intext QuestionIntext Question
Intext Question
7.3 Why is N
2
less reactive at room temperature?
7.3 Ammonia7.3 Ammonia
7.3 Ammonia7.3 Ammonia
7.3 Ammonia
2020-21
176Chemistry
Properties
Ammonia is a colourless gas with a pungent odour. Its freezing and
boiling points are 198.4 and 239.7 K respectively. In the solid and
liquid states, it is associated through hydrogen bonds as in the case
of water and that accounts for its higher melting and boiling points
than expected on the basis of its molecular mass. The ammonia molecule
is trigonal pyramidal with the nitrogen atom at the apex. It has three
bond pairs and one lone pair of electrons as shown in the structure.
Ammonia gas is highly soluble in water. Its aqueous solution is
weakly basic due to the formation of OH
ions.
NH
3
(g) + H
2
O(l) l NH
4
+
(aq) + OH
(aq)
It forms ammonium salts with acids, e.g., NH
4
Cl, (NH
4
)
2
SO
4
, etc. As
a weak base, it precipitates the hydroxides (hydrated oxides in case of
some metals) of many metals from their salt solutions. For example,
(
)
(
)
( ) ( )
(
)
(
)
( )
4 4 4 42
2
ZnSO 2NH OH Zn NH SO
aq aq aq
OH s
white ppt
+ +
(
)
(
)
( )
(
)
(
)
3 4 2 3 2 4
aq aq aq
FeCl NH OH Fe O . H O NH Cl
s
brown ppt
+ +x
H
H
H
N
Fig. 7.1
Flow chart for the
manufacture of
ammonia
On a large scale, ammonia is manufactured by Haber’s process.
N
2
(g) + 3H
2
(g) Ö 2NH
3
(g);
f
H
0
= – 46.1 kJ mol
–1
In accordance with Le Chatelier’s principle, high pressure would
favour the formation of ammonia. The optimum conditions for the
production of ammonia are a pressure of 200 × 10
5
Pa (about 200
atm), a temperature of ~ 700 K and the use of a catalyst such as iron
oxide with small amounts of K
2
O and Al
2
O
3
to increase the rate of
attainment of equilibrium. The flow chart for the production of ammonia
is shown in Fig. 7.1. Earlier, iron was used as a catalyst with
molybdenum as a promoter.
2020-21
177 The p-Block Elements
Example 7.4Example 7.4
Example 7.4Example 7.4
Example 7.4
Why does NH
3
act as a Lewis base ?
Nitrogen atom in NH
3
has one lone pair of electrons which
is available for donation. Therefore, it acts as a Lewis base.
Intext QuestionsIntext Questions
Intext QuestionsIntext Questions
Intext Questions
7.4 Mention the conditions required to maximise the yield of ammonia.
7.5 How does ammonia react with a solution of Cu
2+
?
7.47.4
7.47.4
7.4
Oxides ofOxides of
Oxides ofOxides of
Oxides of
NitrogenNitrogen
NitrogenNitrogen
Nitrogen
SolutionSolution
SolutionSolution
Solution
UsesUses
UsesUses
Uses: Ammonia is used to produce various nitrogenous fertilisers
(ammonium nitrate, urea, ammonium phosphate and ammonium sulphate)
and in the manufacture of some inorganic nitrogen compounds, the most
important one being nitric acid. Liquid ammonia is also used as a refrigerant.
Nitrogen forms a number of oxides in different oxidation states. The
names, formulas, preparation and physical appearance of these oxides
are given in Table 7.3.
Dinitrogen oxide N
2
O + 1
Heat
4 3
2 2
NH NO
N O 2H O

+
colourless gas,
[Nitrogen(I) oxide]
neutral
Nitrogen monoxide NO + 2
2 4 2 4
2NaNO 2FeSO 3H SO
+ +
colourless gas,
[Nitrogen(II) oxide]
(
)
2 4 4
3
Fe SO 2NaHSO
+
neutral
2
2H O 2NO
+ +
Table 7.3: Oxides of Nitrogen
Oxidation
state of
nitrogen
Name Formula Common
methods of
preparation
Physical
appearance and
chemical nature
The presence of a lone pair of electrons on the nitrogen atom of
the ammonia molecule makes it a Lewis base. It donates the electron
pair and forms linkage with metal ions and the formation of such
complex compounds finds applications in detection of metal ions
such as Cu
2+
, Ag
+
:
Cu
2+
(aq) + 4 NH
3
(aq) Ö [Cu(NH
3
)
4
]
2+
(aq)
(blue) (deep blue)
(
)
(
)
( )
Ag Cl AgCl
aq aq
s
+
+
(colourless) (white ppt)
( )
(
)
(
)
(
)
3
3
2
NHAg
AgCl 2NH Cl
aq aq
s
+
(white ppt) (colourless)
2020-21
178Chemistry
Table 7.4: Structures of Oxides of Nitrogen
Lewis dot main resonance structures and bond parameters of oxides
are given in Table 7.4.
Dinitrogen trioxide N
2
O
3
+ 3
250 K
2 4 2 3
2NO N O 2N O
+ 
blue solid,
[Nitrogen(III) oxide] acidic
Nitrogen dioxide NO
2
+ 4
(
)
+ +
673K
3
2
2 2
2Pb NO
4NO 2PbO O
brown gas,
[Nitrogen(IV) oxide] acidic
Dinitrogen tetroxide N
2
O
4
+ 4
Cool
2 2 4
Heat
2NO N O


colourless solid/
liquid, acidic
[Nitrogen(IV) oxide]
Dinitrogen pentoxide N
2
O
5
+5
3 4 10
3 2 5
4HNO P O
4HPO 2N O
+
+
colourless solid,
[Nitrogen(V) oxide] acidic
2020-21
179 The p-Block Elements
Preparation
In the laboratory, nitric acid is prepared by heating KNO
3
or NaNO
3
and concentrated H
2
SO
4
in a glass retort.
3 2 4 4 3
NaNO H SO NaHSO HNO
+ +
On a large scale it is prepared mainly by Ostwald’s process.
This method is based upon catalytic oxidation of NH
3
by atmospheric
oxygen.
(
)
(
)
(
)
(
)
Pt /Rh gauge catalyst
3 2 2
500 K 9 bar,
(from air)
4NH g 5O g 4NO g 6H O g
+  +
Nitric oxide thus formed combines with oxygen giving NO
2
.
(
)
(
)
(
)
2 2
2NO g O g 2NO g
+
Nitrogen dioxide so formed, dissolves in water to give HNO
3
.
(
)
(
)
(
)
(
)
2 2 3
3NO g H O l 2HNO aq NO g
+ +
NO thus formed is recycled and the aqueous HNO
3
can be
concentrated by distillation upto ~ 68% by mass. Further
concentration to 98% can be achieved by dehydration with
concentrated H
2
SO
4
.
Properties
It is a colourless liquid (f.p. 231.4 K and b.p. 355.6 K). Laboratory
grade nitric acid contains ~ 68% of the HNO
3
by mass and has a
specific gravity of 1.504.
In the gaseous state, HNO
3
exists as a planar molecule with
the structure as shown.
In aqueous solution, nitric acid behaves as a strong acid giving
hydronium and nitrate ions.
HNO
3
(aq) + H
2
O(l) H
3
O
+
(aq) + NO
3
(aq)
Concentrated nitric acid is a strong oxidising agent and attacks
most metals except noble metals such as gold and platinum. The
Why does NO
2
dimerise ?
NO
2
contains odd number of valence electrons. It behaves as a typical
odd molecule. On dimerisation, it is converted to stable N
2
O
4
molecule
with even number of electrons.
Intext QuestionIntext Question
Intext QuestionIntext Question
Intext Question
7.6 What is the covalence of nitrogen in N
2
O
5
?
7.5
7.5
7.5
7.5
7.5
Nitric Acid
Nitric Acid
Nitric Acid
Nitric Acid
Nitric Acid
Nitrogen forms oxoacids such as H
2
N
2
O
2
(hyponitrous acid), HNO
2
(nitrous acid) and HNO
3
(nitric acid). Amongst them HNO
3
is the
most important.
SolutionSolution
SolutionSolution
Solution
Example 7.5Example 7.5
Example 7.5Example 7.5
Example 7.5
2020-21
180Chemistry
7.67.6
7.67.6
7.6
Phosphorus —Phosphorus —
Phosphorus —Phosphorus —
Phosphorus —
AllotropicAllotropic
AllotropicAllotropic
Allotropic
FormsForms
FormsForms
Forms
products of oxidation depend upon the concentration of the acid,
temperature and the nature of the material undergoing oxidation.
3Cu + 8 HNO
3
(dilute) 3Cu(NO
3
)
2
+ 2NO + 4H
2
O
Cu + 4HNO
3
(conc.) Cu(NO
3
)
2
+ 2NO
2
+ 2H
2
O
Zinc reacts with dilute nitric acid to give N
2
O and with concentrated
acid to give NO
2
.
4Zn + 10HNO
3
(dilute) 4 Zn (NO
3
)
2
+ 5H
2
O + N
2
O
Zn + 4HNO
3
(conc.) Zn (NO
3
)
2
+ 2H
2
O + 2NO
2
Some metals (e.g., Cr, Al) do not dissolve in concentrated nitric
acid because of the formation of a passive film of oxide on the surface.
Concentrated nitric acid also oxidises non–metals and their
compounds. Iodine is oxidised to iodic acid, carbon to carbon dioxide,
sulphur to H
2
SO
4
, and phosphorus to phosphoric acid.
I
2
+ 10HNO
3
2HIO
3
+ 10NO
2
+ 4H
2
O
C + 4HNO
3
CO
2
+ 2H
2
O + 4NO
2
S
8
+ 48HNO
3
8H
2
SO
4
+ 48NO
2
+ 16H
2
O
P
4
+ 20HNO
3
4H
3
PO
4
+ 20NO
2
+ 4H
2
O
Brown Ring Test: The familiar brown ring test for nitrates depends
on the ability of Fe
2+
to reduce nitrates to nitric oxide, which reacts
with Fe
2+
to form a brown coloured complex. The test is usually carried
out by adding dilute ferrous sulphate solution to an aqueous solution
containing nitrate ion, and then carefully adding concentrated sulphuric
acid along the sides of the test tube. A brown ring at the interface
between the solution and sulphuric acid layers indicates the presence
of nitrate ion in solution.
NO
3
-
+ 3Fe
2+
+ 4H
+
NO + 3Fe
3+
+ 2H
2
O
+2
62
])OH( Fe[
+ NO [Fe (H
2
O)
5
(NO)]
2+
+ H
2
O
(brown)
UsesUses
UsesUses
Uses: The major use of nitric acid is in the manufacture of ammonium nitrate
for fertilisers and other nitrates for use in explosives and pyrotechnics. It is
also used for the preparation of nitroglycerin, trinitrotoluene and other organic
nitro compounds. Other major uses are in the pickling of stainless steel,
etching of metals and as an oxidiser in rocket fuels.
Phosphorus is found in many allotropic forms, the important ones
being white, red and black.
White phosphorus is a translucent white waxy solid. It is poisonous,
insoluble in water but soluble in carbon disulphide and glows in dark
(chemiluminescence). It dissolves in boiling NaOH solution in an inert
atmosphere giving PH
3
.
(
)
4 2 3 2 2
sodium hypophosphite
P 3NaOH 3H O PH 3NaH PO+ + +
2020-21
181 The p-Block Elements
White phosphorus is less stable and therefore, more reactive than
the other solid phases under normal conditions because of angular
strain in the P
4
molecule where the angles are only 60°. It readily
catches fire in air to give dense white fumes of P
4
O
10
.
4 2 4 10
P 5O P O
+
It consists of discrete tetrahedral P
4
molecule as shown in Fig. 7.2.
Red phosphorus is obtained by heating white phosphorus at 573K
in an inert atmosphere for several days. When red phosphorus is heated
under high pressure, a series of phases of black phosphorus is formed.
Red phosphorus possesses iron grey lustre. It is odourless, non-
poisonous and insoluble in water as well as in carbon disulphide.
Chemically, red phosphorus is much less reactive than white
phosphorus. It does not glow in the dark.
It is polymeric, consisting of chains of P
4
tetrahedra linked together in the manner as shown
in Fig. 7.3.
Black phosphorus has two forms α-black
phosphorus and β-black phosphorus. α-Black
phosphorus is formed when red phosphorus is
heated in a sealed tube at 803K. It can be sublimed
in air and has opaque monoclinic or rhombohedral
crystals. It does not oxidise in air. β-Black phosphorus is prepared by
heating white phosphorus at 473 K under high pressure. It does not
burn in air upto 673 K.
Preparation
Phosphine is prepared by the reaction of calcium phosphide with water
or dilute HCl.
Ca
3
P
2
+ 6H
2
O 3Ca(OH)
2
+ 2PH
3
Ca
3
P
2
+ 6HCl 3CaCl
2
+ 2PH
3
In the laboratory, it is prepared by heating white phosphorus with
concentrated NaOH solution in an inert atmosphere of CO
2
.
(
)
4 2 3 2 2
P 3NaOH 3H O PH 3NaH PO
sodium hypophosphite
+ + +
When pure, it is non inflammable but becomes inflammable owing
to the presence of P
2
H
4
or P
4
vapours. To purify it from the impurities,
it is absorbed in HI to form phosphonium iodide (PH
4
I) which on treating
with KOH gives off phosphine.
4 2 3
PH I KOH KI H O PH
+ + +
Properties
It is a colourless gas with rotten fish smell and is highly poisonous.
It explodes in contact with traces of oxidising agents like HNO
3
, Cl
2
and
Br
2
vapours.
It is slightly soluble in water. The solution of PH
3
in water decomposes
in presence of light giving red phosphorus and H
2
. When absorbed in
P P
P
P P
P
P P
P
P
P
P
Fig.7.3: Red phosphorus
P
P P
P
60°
Fig. 7.2
White phosphorus
7.7
7.7
7.7
7.7
7.7
Phosphine
Phosphine
Phosphine
Phosphine
Phosphine
2020-21
182Chemistry
7.8.1 Phosphorus
Trichloride
Phosphorus forms two types of halides, PX
3
(X = F, Cl, Br, I) and
PX
5
(X = F, Cl, Br).
Preparation
It is obtained by passing dry chlorine over heated white phosphorus.
4 2 3
P 6Cl 4PCl
+
It is also obtained by the action of thionyl chloride with white
phosphorus.
4 2 3 2 2 2
P 8SOCl 4PCl 4SO 2S Cl
+ + +
Properties
It is a colourless oily liquid and hydrolyses in the presence of moisture.
3 2 3 3
PCl 3H O H PO + 3HCl
+
It reacts with organic compounds containing –OH group such as
CH
3
COOH, C
2
H
5
OH.
3 3 3 3 3
3CH COOH PCl 3CH COCl H PO
+ +
2 5 3 2 5 3 3
3C H OH PCl 3C H Cl H PO
+ +
It has a pyramidal shape as shown, in which phosphorus is sp
3
hybridised.
7.87.8
7.87.8
7.8
Phosphorus HalidesPhosphorus Halides
Phosphorus HalidesPhosphorus Halides
Phosphorus Halides
UsesUses
Uses
Uses
Uses: The spontaneous combustion of phosphine is technically used in Holme’s
signals. Containers containing calcium carbide and calcium phosphide are pierced
and thrown in the sea when the gases evolved burn and serve as a signal. It is also
used in smoke screens.
P
Cl
Cl
Cl
7.8.2 Phosphorus
Pentachloride
copper sulphate or mercuric chloride solution, the corresponding
phosphides are obtained.
4 3 3 2 2 4
3CuSO 2PH Cu P 3H SO
+ +
2 3 3 2
3HgCl 2PH Hg P 6HCl
+ +
Phosphine is weakly basic and like ammonia, gives phosphonium
compounds with acids e.g.,
3 4
PH HBr PH Br
+
Example 7.6Example 7.6
Example 7.6Example 7.6
Example 7.6
In what way can it be proved that PH
3
is basic in nature?
PH
3
reacts with acids like HI to form PH
4
I which shows that it is
basic in nature.
3 4
PH HI PH I
+
Due to lone pair on phosphorus atom, PH
3
is acting as a Lewis base in the above reaction.
SolutionSolution
SolutionSolution
Solution
Intext QuestionsIntext Questions
Intext QuestionsIntext Questions
Intext Questions
7.7 (a) Bond angle in PH
4
+
is higher than that in PH
3
. Why?
(b) What is formed when PH
3
reacts with an acid?
7.8 What happens when white phosphorus is heated with concentrated NaOH
solution in an inert atmosphere of CO
2
?
2020-21
183 The p-Block Elements
Preparation
Phosphorus pentachloride is prepared by the reaction of white
phosphorus with excess of dry chlorine.
4 2 5
P 10Cl 4PCl
+
It can also be prepared by the action of SO
2
Cl
2
on phosphorus.
4 2 2 5 2
P 10SO Cl 4PCl 10SO
+ +
Properties
PCl
5
is a yellowish white powder and in moist air, it hydrolyses to
POCl
3
and finally gets converted to phosphoric acid.
5 2 3
PCl H O POCl 2HCl
+ +
3 2 3 4
POCl 3H O H PO 3HCl
+ +
When heated, it sublimes but decomposes on stronger heating.
Heat
5 3 2
PCl PCl Cl
+
It reacts with organic compounds containing –OH group converting
them to chloro derivatives.
2 5 5 2 5 3
C H OH PCl C H Cl POCl HCl
+ + +
3 5 3 3
CH COOH PCl CH COCl POCl +HCl
+ +
Finely divided metals on heating with PCl
5
give corresponding
chlorides.
5 3
5 4 3
2Ag PCl 2AgCl PCl
Sn 2PCl SnCl 2PCl
+ +
+ +
It is used in the synthesis of some organic
compounds, e.g., C
2
H
5
Cl, CH
3
COCl.
In gaseous and liquid phases, it has a trigonal
bipyramidal structure as shown. The three equatorial
P–Cl bonds are equivalent, while the two axial bonds are
longer than equatorial bonds. This is due to the fact that
the axial bond pairs suffer more repulsion as compared
to equatorial bond pairs.
P
Cl
Cl
Cl
Cl
Cl
240 pm
202 pm
Why does PCl
3
fume in moisture ?
PCl
3
hydrolyses in the presence of moisture giving fumes of HCl.
3 2 3 3
PCl 3H O H PO + 3HCl
+
Are all the five bonds in PCl
5
molecule equivalent? Justify your answer.
PCl
5
has a trigonal bipyramidal structure and the three equatorial
P-Cl bonds are equivalent, while the two axial bonds are different and
longer than equatorial bonds.
Example 7.7Example 7.7
Example 7.7Example 7.7
Example 7.7
SolutionSolution
SolutionSolution
Solution
Example 7.8Example 7.8
Example 7.8Example 7.8
Example 7.8
SolutionSolution
SolutionSolution
Solution
2020-21
184Chemistry
Intext QuestionsIntext Questions
Intext QuestionsIntext Questions
Intext Questions
7.9 What happens when PCl
5
is heated?
7.10 Write a balanced equation for the reaction of PCl
5
with water.
Phosphorus forms a number of oxoacids. The important oxoacids of
phosphorus with their formulas, methods of preparation and the
presence of some characteristic bonds in their structures are given
in Table 7.5.
* Exists in polymeric forms only. Characteristic bonds of (HPO
3
)
3
have been given in the Table.
The compositions of the oxoacids are interrelated in terms of loss
or gain of H
2
O molecule or O-atom. The structures of some important
oxoacids are given next.
In oxoacids phosphorus is tetrahedrally surrounded by other atoms.
All these acids contain at least one P=O bond and one P–OH bond. The
oxoacids in which phosphorus has lower oxidation state (less than +5)
contain, in addition to P=O and P–OH bonds, either P–P (e.g., in H
4
P
2
O
6
)
or P–H (e.g., in H
3
PO
2
) bonds but not both. These acids in +3 oxidation
state of phosphorus tend to disproportionate to higher and lower
oxidation states. For example, orthophophorous acid (or phosphorous
acid) on heating disproportionates to give orthophosphoric acid (or
phosphoric acid) and phosphine.
3 3 3 4 3
4H PO 3H PO PH
+
Table 7.5: Oxoacids of Phosphorus
7.9
7.9
7.9
7.9
7.9
Oxoacids of
Oxoacids of
Oxoacids of
Oxoacids of
Oxoacids of
Phosphorus
Phosphorus
Phosphorus
Phosphorus
Phosphorus
Hypophosphorous H
3
PO
2
+1 One P – OH white P
4
+ alkali
(Phosphinic) Two P – H
One P = O
Orthophosphorous H
3
PO
3
+3 Two P – OH P
2
O
3
+ H
2
O
(Phosphonic) One P – H
One P = O
Pyrophosphorous H
4
P
2
O
5
+3 Two P – OH PCl
3
+ H
3
PO
3
Two P – H
Two P = O
Hypophosphoric H
4
P
2
O
6
+4 Four P – OH red P
4
+ alkali
Two P = O
One P – P
Orthophosphoric H
3
PO
4
+5 Three P – OH P
4
O
10
+H
2
O
One P = O
Pyrophosphoric H
4
P
2
O
7
+5 Four P – OH heat phosphoric
Two P = O acid
One P – O – P
Metaphosphoric* (HPO
3
)
n
+5 Three P – OH phosphorus acid
Three P = O + Br
2
, heat in a
Three P – O – P sealed tube
Oxidation
state of
phosphorus
Characteristic
bonds and their
number
PreparationName Formula
2020-21
185 The p-Block Elements
The acids which contain P–H bond have strong reducing properties.
Thus, hypophosphorous acid is a good reducing agent as it contains
two P–H bonds and reduces, for example, AgNO
3
to metallic silver.
4 AgNO
3
+ 2H
2
O + H
3
PO
2
4Ag + 4HNO
3
+ H
3
PO
4
These P–H bonds are not ionisable to give H
+
and do not play any
role in basicity. Only those H atoms which are attached with oxygen in
P–OH form are ionisable and cause the basicity. Thus, H
3
PO
3
and
H
3
PO
4
are dibasic and tribasic, respectively as the structure of H
3
PO
3
has two P–OH bonds and H
3
PO
4
three.
Oxygen, sulphur, selenium, tellurium, polonium and livermorium
constitute Group 16 of the periodic table. This is sometimes known as
group of chalcogens. The name is derived from the Greek word for
brass and points to the association of sulphur and its congeners with
copper. Most copper minerals contain either oxygen or sulphur and
frequently the other members of the group.
7.107.10
7.107.10
7.10
Group 16Group 16
Group 16Group 16
Group 16
ElementsElements
ElementsElements
Elements
Fig. 7.4
Structures of some
important oxoacids of
phosphorus
O
P
HO
H PO
Orthophosphoric acid
3 4
OH
OH
OH
OH
O
P
Orthophosphorous acid
H PO
3 3
H
HO
O
O
P
P
O
H P O
Pyrophosphoric acid
4 2 7
OH
OH
OH
O
P
H
H
OH
Hypophosphorous acid
H PO
3 2
O
O
O
O
O
O
HO
OH
OH
P
P
P
O
O
O
O
O O
OH OH
P
P
P
OH O
Cyclotrimetaphosphoric acid, (HPO )
3 3
Polymetaphosphoric acid, (HPO )
3 n
Intext QuestionsIntext Questions
Intext QuestionsIntext Questions
Intext Questions
7.11 What is the basicity of H
3
PO
4
?
7.12 What happens when H
3
PO
3
is heated?
How do you account for the reducing behaviour
of H
3
PO
2
on the basis of its structure ?
In H
3
PO
2
, two H atoms are bonded directly to P
atom which imparts reducing character to the acid.
Example 7.9Example 7.9
Example 7.9Example 7.9
Example 7.9
SolutionSolution
SolutionSolution
Solution
2020-21
186Chemistry
Oxygen is the most abundant of all the elements on earth. Oxygen
forms about 46.6% by mass of earth’s crust. Dry air contains 20.946%
oxygen by volume.
However, the abundance of sulphur in the earth’s crust is only
0.03-0.1%. Combined sulphur exists primarily as sulphates such as
gypsum CaSO
4
.2H
2
O, epsom salt MgSO
4
.7H
2
O, baryte BaSO
4
and
sulphides such as galena PbS, zinc blende ZnS, copper pyrites CuFeS
2
.
Traces of sulphur occur as hydrogen sulphide in volcanoes. Organic
materials such as eggs, proteins, garlic, onion, mustard, hair and wool
contain sulphur.
Selenium and tellurium are also found as metal selenides and
tellurides in sulphide ores. Polonium occurs in nature as a decay
product of thorium and uranium minerals. Livermorium is a synthetic
radioactive element. Its symbol is Lv, atomic number 116, atomic mass
292 and electronic configuration [Rn] 5f
14
6d
10
7s
2
7p
4
. It has been
produced only in a very small amount and has very short half-life (only
a small fraction of one second). This limits the study of properlies of Lv.
Here, except for livermorium, important atomic and physical
properties of other elements of Group16 along with their electronic
configurations are given in Table 7.6. Some of the atomic, physical and
chemical properties and their trends are discussed below.
7.10.2 Electronic
Configuration
7.10.3 Atomic
and Ionic
Radii
The elements of Group16 have six electrons in the outermost shell and
have ns
2
np
4
general electronic configuration.
Due to increase in the number of shells, atomic and ionic radii increase
from top to bottom in the group. The size of oxygen atom is, however,
exceptionally small.
a
Single bond;
b
Approximate value;
c
At the melting point;
d
Rhombic sulphur;
e
Hexagonal grey;
f
Monoclinic form, 673 K.
* Oxygen shows oxidation states of +2 and +1 in oxygen fluorides OF
2
and O
2
F
2
respectively.
Atomic number 8 16 34 52 84
Atomic mass/g mol
–1
16.00 32.06 78.96 127.60 210.00
Electronic configuration [He]2s
2
2p
4
[Ne]3s
2
3p
4
[Ar]3d
10
4s
2
4p
4
[Kr]4d
10
5s
2
5p
4
[Xe]4f
14
5d
10
6s
2
6p
4
Covalent radius/(pm)
a
66 104 117 137 146
Ionic radius, E
2–
/pm 140 184 198 221 230
b
Electron gain enthalpy, –141 –200 –195 –190 –174
/
eg
H kJ mol
–1
Ionisation enthalpy (
i
H
1
) 1314 1000 941 869 813
/kJ mol
–1
Electronegativity 3.50
2.58 2.55 2.01 1.76
Density /g cm
–3
(298 K) 1.32
c
2.06
d
4.19
e
6.25
Melting point/K 55 393
f
490 725 520
Boiling point/K 90 718 958 1260 1235
Oxidation states* –2,–1,1,2 –2,2,4,6 –2,2,4,6 –2,2,4,6 2,4
Property O S Se Te Po
Table 7.6: Some Physical Properties of Group 16 Elements
7.10.1 Occurrence
2020-21
187 The p-Block Elements
7.10.4 Ionisation
Enthalpy
7.10.5 Electron
Gain
Enthalpy
7.10.6
Electronegativity
Ionisation enthalpy decreases down the group. It is due to increase in
size. However, the elements of this group have lower ionisation enthalpy
values compared to those of Group15 in the corresponding periods.
This is due to the fact that Group 15 elements have extra stable half-
filled p orbitals electronic configurations.
Because of the compact nature of oxygen atom, it has less negative
electron gain enthalpy than sulphur. However, from sulphur onwards
the value again becomes less negative upto polonium.
Next to fluorine, oxygen has the highest electronegativity value amongst
the elements. Within the group, electronegativity decreases with an
increase in atomic number. This implies that the metallic character
increases from oxygen to polonium.
Some of the physical properties of Group 16 elements are given in
Table 7.6. Oxygen and sulphur are non-metals, selenium and tellurium
metalloids, whereas polonium is a metal. Polonium is radioactive and
is short lived (Half-life 13.8 days). All these elements exhibit allotropy.
The melting and boiling points increase with an increase in atomic
number down the group. The large difference between the melting and
boiling points of oxygen and sulphur may be explained on the basis
of their atomicity; oxygen exists as diatomic molecule (O
2
) whereas
sulphur exists as polyatomic molecule (S
8
).
Oxidation states and trends in chemical reactivity
The elements of Group 16 exhibit a number of oxidation states (Table
7.6). The stability of -2 oxidation state decreases down the group.
Polonium hardly shows –2 oxidation state. Since electronegativity of
oxygen is very high, it shows only negative oxidation state as –2 except
in the case of OF
2
where its oxidation state is + 2. Other elements of the
group exhibit + 2, + 4, + 6 oxidation states but + 4 and + 6 are more
common. Sulphur, selenium and tellurium usually show + 4 oxidation
state in their compounds with oxygen and + 6 with fluorine. The stability
of + 6 oxidation state decreases down the group and stability of + 4
oxidation state increases (inert pair effect). Bonding in +4 and +6
oxidation states is primarily covalent.
Anomalous behaviour of oxygen
The anomalous behaviour of oxygen, like other members of p-block
present in second period is due to its small size and high
electronegativity. One typical example of effects of small size and high
electronegativity is the presence of strong hydrogen bonding in H
2
O
which is not found in H
2
S.
7.10.7 Physical
Properties
7.10.8 Chemical
Properties
Elements of Group 16 generally show lower value of first ionisation
enthalpy compared to the corresponding periods of group 15. Why?
Due to extra stable half-filled p orbitals electronic configurations of
Group 15 elements, larger amount of energy is required to remove
electrons compared to Group 16 elements.
Example 7.10Example 7.10
Example 7.10Example 7.10
Example 7.10
SolutionSolution
SolutionSolution
Solution
2020-21
188Chemistry
(ii) Reactivity with oxygen: All these elements form oxides of the EO
2
and EO
3
types where E = S, Se, Te or Po. Ozone (O
3
) and sulphur
dioxide (SO
2
) are gases while selenium dioxide (SeO
2
) is solid.
Reducing property of dioxide decreases from SO
2
to TeO
2
; SO
2
is
reducing while TeO
2
is an oxidising agent. Besides EO
2
type,
sulphur, selenium and tellurium also form EO
3
type oxides (SO
3
,
SeO
3
, TeO
3
). Both types of oxides are acidic in nature.
(iii) Reactivity towards the halogens: Elements of Group 16 form a
large number of halides of the type, EX
6
, EX
4
and EX
2
where E is
an element of the group and X is a halogen. The stability of the
halides decreases in the order F
> Cl
> Br
> I
. Amongst hexahalides,
hexafluorides are the only stable halides. All hexafluorides are
gaseous in nature. They have octahedral structure. Sulphur
hexafluoride, SF
6
is exceptionally stable for steric reasons.
Amongst tetrafluorides, SF
4
is a gas, SeF
4
a liquid and TeF
4
a solid.
These fluorides have sp
3
d hybridisation and thus, have trigonal
bipyramidal structures in which one of the equatorial positions is
occupied by a lone pair of electrons. This geometry is also regarded as
see-saw geometry.
All elements except oxygen form dichlorides and dibromides. These
dihalides are formed by sp
3
hybridisation and thus, have tetrahedral
structure. The well known monohalides are dimeric in nature. Examples
The absence of d orbitals in oxygen limits its covalency to four and
in practice, rarely exceeds two. On the other hand, in case of other
elements of the group, the valence shells can be expanded and covalence
exceeds four.
(i) Reactivity with hydrogen: All the elements of Group 16 form
hydrides of the type H
2
E (E = O, S, Se, Te, Po). Some properties of
hydrides are given in Table 7.7. Their acidic character increases
from H
2
O to H
2
Te. The increase in acidic character can be explained
in terms of decrease in bond enthalpy for the dissociation of H–E
bond down the group. Owing to the decrease in enthalpy for the
dissociation of H–E bond down the group, the thermal stability of
hydrides also decreases from H
2
O to H
2
Po. All the hydrides except
water possess reducing property and this character increases from
H
2
S to H
2
Te.
Property H
2
O H
2
S H
2
Se H
2
Te
m.p/K 273 188 208 222
b.p/K 373 213 232 269
H–E distance/pm 96 134 146 169
HEH angle (°) 104 92 91 90
f
H/kJ mol
–1
–286 –20 73 100
diss
H (H–E)/kJ mol
–1
463 347 276
238
Dissociation constant
a
1.8×10
–16
1.3×10
–7
1.3×10
–4
2.3×10
–3
a
Aqueous solution, 298 K
Table 7.7: Properties of Hydrides of Group 16 Elements
2020-21
189 The p-Block Elements
are S
2
F
2
, S
2
Cl
2
, S
2
Br
2
, Se
2
Cl
2
and Se
2
Br
2
. These dimeric halides undergo
disproportionation as given below:
2Se
2
Cl
2
SeCl
4
+ 3Se
7.117.11
7.117.11
7.11
DioxygenDioxygen
DioxygenDioxygen
Dioxygen
Preparation
Dioxygen can be obtained in the laboratory by the following ways:
(i) By heating oxygen containing salts such as chlorates, nitrates and
permanganates.
2
Heat
3 2
MnO
2KClO 2KCl 3O
+
(ii) By the thermal decomposition of the oxides of metals low in the
electrochemical series and higher oxides of some metals.
2Ag
2
O(s) 4Ag(s) + O
2
(g); 2Pb
3
O
4
(s) 6PbO(s) + O
2
(g)
2HgO(s) 2Hg(l) + O
2
(g) ; 2PbO
2
(s) 2PbO(s) + O
2
(g)
(iii) Hydrogen peroxide is readily decomposed into water and dioxygen
by catalysts such as finely divided metals and manganese dioxide.
2H
2
O
2
(aq) 2H
2
O(1) + O
2
(g)
On large scale it can be prepared from water or air. Electrolysis of
water leads to the release of hydrogen at the cathode and oxygen
at the anode.
Industrially, dioxygen is obtained from air by first removing carbon
dioxide and water vapour and then, the remaining gases are liquefied
and fractionally distilled to give dinitrogen and dioxygen.
Properties
Dioxygen is a colourless and odourless gas. Its solubility in water is to
the extent of 3.08 cm
3
in 100 cm
3
water at 293 K which is just sufficient
for the vital support of marine and aquatic life. It liquefies at 90 K and
freezes at 55 K. Oxygen atom has three stable isotopes:
16
O,
17
O and
18
O. Molecular oxygen, O
2
is unique in being paramagnetic inspite of
having even number of electrons (see Class XI Chemistry Book, Unit 4).
Dioxygen directly reacts with nearly all metals and non-metals
except some metals ( e.g., Au, Pt) and some noble gases. Its combination
with other elements is often strongly exothermic which helps in
sustaining the reaction. However, to initiate the reaction, some external
H
2
S is less acidic than H
2
Te. Why?
Due to the decrease in bond (E–H) dissociation
enthalpy down the group, acidic character increases.
Example 7.11Example 7.11
Example 7.11
Example 7.11
Example 7.11
SolutionSolution
Solution
Solution
Solution
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7.13 List the important sources of sulphur.
7.14 Write the order of thermal stability of the hydrides of Group 16 elements.
7.15 Why is H
2
O a liquid and H
2
S a gas ?
2020-21
190Chemistry
A binary compound of oxygen with another element is called oxide. As
already stated, oxygen reacts with most of the elements of the periodic
table to form oxides. In many cases one element forms two or more
oxides. The oxides vary widely in their nature and properties.
Oxides can be simple (e.g., MgO, Al
2
O
3
) or mixed (Pb
3
O
4
, Fe
3
O
4
).
Simple oxides can be classified on the basis of their acidic, basic or
amphoteric character. An oxide that combines with water to give an
acid is termed acidic oxide (e.g., SO
2
, Cl
2
O
7
, CO
2
, N
2
O
5
). For example,
SO
2
combines with water to give H
2
SO
3
, an acid.
2 2 2 3
SO H O H SO
+
As a general rule, only non-metal oxides are acidic but oxides of
some metals in high oxidation state also have acidic character (e.g.,
Mn
2
O
7
, CrO
3
, V
2
O
5
). The oxides which give a base with water are known
as basic oxides (e.g., Na
2
O, CaO, BaO). For example, CaO combines
with water to give Ca(OH)
2
, a base.
(
)
2
2
CaO H O Ca
OH
+
7.16 Which of the following does not react with oxygen directly?
Zn, Ti, Pt, Fe
7.17 Complete the following reactions:
(i) C
2
H
4
+ O
2
(ii) 4Al + 3 O
2
7.127.12
7.127.12
7.12
SimpleSimple
SimpleSimple
Simple
OxidesOxides
OxidesOxides
Oxides
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heating is required as bond dissociation enthalpy of oxgyen-oxygen
double bond is high (493.4 kJ mol
–1
).
Some of the reactions of dioxygen with metals, non-metals and
other compounds are given below:
2
2Ca O 2CaO
+
2 2 3
4Al 3O 2Al O
+
4 2 4 10
P 5O P O
+
2 2
C O CO
+
2ZnS + 3O
2
2ZnO + 2SO
2
4 2 2 2
CH 2O CO 2H O
+ +
Some compounds are catalytically oxidised. For example,
2 5
V O
2 2 3
2SO O 2SO
+
2
CuCl
2 2 2
4HCl O 2Cl 2H O
+  +
UsesUses
UsesUses
Uses: In addition to its importance in normal respiration and combustion
processes, oxygen is used in oxyacetylene welding, in the manufacture of
many metals, particularly steel. Oxygen cylinders are widely used in hospitals,
high altitude flying and in mountaineering. The combustion of fuels, e.g.,
hydrazines in liquid oxygen, provides tremendous thrust in rockets.
2020-21
191 The p-Block Elements
In general, metallic oxides are basic.
Some metallic oxides exhibit a dual behaviour. They show
characteristics of both acidic as well as basic oxides. Such oxides are
known as amphoteric oxides. They react with acids as well as alkalies.
For example, Al
2
O
3
reacts with acids as well as alkalies.
( )
( )
( )
[
]
( ) ( )
( )
( )
( )
( )
[ ]
( )
3
2 62 3 2
2 3 2 3 6
Al(H O)Al O 6HCl 9H O 2 6Cl
aq aq aq
s
l
Al
Al O 6NaOH 3H O 2Naaq aq
OH
s l
+
+ + +
+ +
There are some oxides which are neither acidic nor basic. Such oxides
are known as neutral oxides. Examples of neutral oxides are CO, NO
and N
2
O.
Ozone is an allotropic form of oxygen. It is too reactive to remain for
long in the atmosphere at sea level. At a height of about 20 kilometres,
it is formed from atmospheric oxygen in the presence of sunlight. This
ozone layer protects the earth’s surface from an excessive concentration
of ultraviolet (UV) radiations.
Preparation
When a slow dry stream of oxygen is passed through a silent electrical
discharge, conversion of oxygen to ozone (10%) occurs. The product is
known as ozonised oxygen.
3O
2
2O
3
H
V
(298 K) = +142 kJ mol
–1
Since the formation of ozone from oxygen is an endothermic process,
it is necessary to use a silent electrical discharge in its preparation to
prevent its decomposition.
If concentrations of ozone greater than 10 per cent are required, a
battery of ozonisers can be used, and pure ozone (b.p. 101.1K) can be
condensed in a vessel surrounded by liquid oxygen.
Properties
Pure ozone is a pale blue gas, dark blue liquid and violet-black solid.
Ozone has a characteristic smell and in small concentrations it is harmless.
However, if the concentration rises above about 100 parts per million,
breathing becomes uncomfortable resulting in headache and nausea.
Ozone is thermodynamically unstable with respect to oxygen since
its decomposition into oxygen results in the liberation of heat (
H is
negative) and an increase in entropy (
S is positive). These two effects
reinforce each other, resulting in large negative Gibbs energy change
(G) for its conversion into oxygen. It is not really surprising, therefore,
high concentrations of ozone can be dangerously explosive.
Due to the ease with which it liberates atoms of nascent oxygen
(O
3
O
2
+ O), it acts as a powerful oxidising agent. For example, it
oxidises lead sulphide to lead sulphate and iodide ions to iodine.
PbS(s) + 4O
3
(g) PbSO
4
(s) + 4O
2
(g)
2I
(aq) + H
2
O(l) + O
3
(g) 2OH
(aq) + I
2
(s) + O
2
(g)
When ozone reacts with an excess of potassium iodide solution
buffered with a borate buffer (pH 9.2), iodine is liberated which can be
titrated against a standard solution of sodium thiosulphate. This is a
quantitative method for estimating O
3
gas.
7.13
7.13
7.13
7.13
7.13
Ozone
Ozone
Ozone
Ozone
Ozone
2020-21
192Chemistry
Sulphur forms numerous allotropes of which the yellow rhombic
(α-sulphur) and monoclinic (β -sulphur) forms are the most important.
The stable form at room temperature is rhombic sulphur, which
transforms to monoclinic sulphur when heated above 369 K.
Rhombic sulphur (
αα
αα
α-sulphur)
This allotrope is yellow in colour, m.p. 385.8 K and specific gravity
2.06. Rhombic sulphur crystals are formed on evaporating the solution
of roll sulphur in CS
2
. It is insoluble in water but dissolves to some
extent in benzene, alcohol and ether. It is readily soluble in CS
2
.
Monoclinic sulphur (
β
β
ββ
β-sulphur)
Its m.p. is 393 K and specific gravity 1.98. It is soluble in CS
2
. This
form of sulphur is prepared by melting rhombic sulphur in a dish
and cooling, till crust is formed. Two holes are made in the crust and
the remaining liquid poured out. On removing the crust, colourless
needle shaped crystals of β-sulphur are formed. It is stable above 369 K
and transforms into α-sulphur below it. Conversely, α-sulphur is stable
below 369 K and transforms into β-sulphur above this. At 369 K both
the forms are stable. This temperature is called transition temperature.
Both rhombic and monoclinic sulphur have S
8
molecules. These S
8
molecules are packed to give different crystal structures. The S
8
ring
in both the forms is puckered and has a crown shape. The molecular
dimensions are given in Fig. 7.5(a).
7.14 Sulphur —7.14 Sulphur —
7.14 Sulphur —7.14 Sulphur —
7.14 Sulphur —
AllotropicAllotropic
AllotropicAllotropic
Allotropic
FormsForms
FormsForms
Forms
Experiments have shown that nitrogen oxides (particularly nitrogen
monoxide) combine very rapidly with ozone and there is, thus, the
possibility that nitrogen oxides emitted from the exhaust systems of
supersonic jet aeroplanes might be slowly depleting the concentration
of the ozone layer in the upper atmosphere.
(
)
(
)
(
)
(
)
3 2 2
NO g O g NO g O g
+ +
Another threat to this ozone layer is probably posed by the use of
freons which are used in aerosol sprays and as refrigerants.
The two oxygen-oxygen bond lengths in the ozone
molecule are identical (128 pm) and the molecule is angular
as expected with a bond angle of about 117
o
. It is a resonance
hybrid of two main forms:
7.18 Why does O
3
act as a powerful oxidising agent?
7.19 How is O
3
estimated quantitatively?
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UsesUses
UsesUses
Uses: It is used as a germicide, disinfectant and for sterilising water. It is also
used for bleaching oils, ivory, flour, starch, etc. It acts as an oxidising agent
in the manufacture of potassium permanganate.
2020-21
193 The p-Block Elements
Which form of sulphur shows paramagnetic behaviour ?
In vapour state sulphur partly exists as S
2
molecule which has two
unpaired electrons in the antibonding
π
* orbitals like O
2
and, hence,
exhibits paramagnetism.
Example 7.12Example 7.12
Example 7.12Example 7.12
Example 7.12
SolutionSolution
SolutionSolution
Solution
Several other modifications
of sulphur containing 6-20
sulphur atoms per ring have
been synthesised in the last
two decades. In cyclo-S
6
, the
ring adopts the chair form and
the molecular dimensions are
as shown in Fig. 7.5 (b).
At elevated temperatures
(~1000 K), S
2
is the dominant
species and is paramagnetic
like O
2
.
Preparation
Sulphur dioxide is formed together with a little (6-8%) sulphur trioxide
when sulphur is burnt in air or oxygen:
S(s) + O
2
(g) SO
2
(g)
In the laboratory it is readily generated by treating a sulphite with
dilute sulphuric acid.
SO
3
2-
(aq) + 2H
+
(aq) H
2
O(l) + SO
2
(g)
Industrially, it is produced as a by-product of the roasting of
sulphide ores.
( )
(
)
( )
(
)
2 2 2 3 2
4FeS 11O 2Fe O 8SO
g g
s s
+ +
The gas after drying is liquefied under pressure and stored in steel cylinders.
Properties
Sulphur dioxide is a colourless gas with pungent smell and is highly
soluble in water. It liquefies at room temperature under a pressure of
two atmospheres and boils at 263 K.
Sulphur dioxide, when passed through water, forms a solution of
sulphurous acid.
SO g H O l H SO
2 2 2 3
( )
+
( ) ( )
aq
It reacts readily with sodium hydroxide solution, forming sodium
sulphite, which then reacts with more sulphur dioxide to form sodium
hydrogen sulphite.
2NaOH + SO
2
Na
2
SO
3
+ H
2
O
Na
2
SO
3
+ H
2
O + SO
2
2NaHSO
3
In its reaction with water and alkalies, the behaviour of sulphur
dioxide is very similar to that of carbon dioxide.
7.157.15
7.157.15
7.15
SulphurSulphur
SulphurSulphur
Sulphur
DioxideDioxide
DioxideDioxide
Dioxide
(a) (b)
Fig. 7.5: The structures of (a) S
8
ring in
rhombic sulphur and (b) S
6
form
2020-21
194Chemistry
Sulphur forms a number of oxoacids such as H
2
SO
3
, H
2
S
2
O
3
, H
2
S
2
O
4
,
H
2
S
2
O
5
, H
2
S
x
O
6
(x = 2 to 5), H
2
SO
4
, H
2
S
2
O
7
, H
2
SO
5
, H
2
S
2
O
8
. Some of
these acids are unstable and cannot be isolated. They are known in
aqueous solution or in the form of their salts. Structures of some
important oxoacids are shown in Fig. 7.6.
7.167.16
7.167.16
7.16
Oxoacids ofOxoacids of
Oxoacids ofOxoacids of
Oxoacids of
SulphurSulphur
SulphurSulphur
Sulphur
Fig. 7.6: Structures of some important oxoacids of sulphur
Sulphur dioxide reacts with chlorine in the presence of charcoal (which
acts as a catalyst) to give sulphuryl chloride, SO
2
Cl
2
. It is oxidised to
sulphur trioxide by oxygen in the presence of vanadium(V) oxide catalyst.
SO
2
(g) + Cl
2
(g) SO
2
Cl
2
(l)
(
)
(
)
(
)
2 5
V O
2 2 3
2SO g O g 2SO g
+
When moist, sulphur dioxide behaves as a reducing agent. For
example, it converts iron(III) ions to iron(II) ions and decolourises
acidified potassium permanganate(VII) solution; the latter reaction is a
convenient test for the gas.
3 2 2
2 2 4
2 2
2 4 2 4
2Fe SO 2H O 2Fe SO 4H
5SO 2MnO 2H O 5SO 4H 2Mn
+ + +
+ +
+ + + +
+ + + +
The molecule of SO
2
is angular. It is a resonance hybrid
of the two canonical forms:
Uses
Uses
Uses
Uses
Uses: Sulphur dioxide is used (i) in refining petroleum and sugar (ii) in bleaching
wool and silk and (iii) as an anti-chlor, disinfectant and preservative. Sulphuric
acid, sodium hydrogen sulphite and calcium hydrogen sulphite (industrial
chemicals) are manufactured from sulphur dioxide. Liquid SO
2
is used as a
solvent to dissolve a number of organic and inorganic chemicals.
Intext QuestionsIntext Questions
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Intext Questions
7.20 What happens when sulphur dioxide is passed through an aqueous
solution of Fe(III) salt?
7.21 Comment on the nature of two S–O bonds formed in SO
2
molecule. Are
the two S–O bonds in this molecule equal ?
7.22 How is the presence of SO
2
detected ?
2020-21
195 The p-Block Elements
Manufacture
Sulphuric acid is one of the most important industrial chemicals
worldwide.
Sulphuric acid is manufactured by the Contact Process which involves
three steps:
(i) burning of sulphur or sulphide ores in air to generate SO
2
.
(ii) conversion of SO
2
to SO
3
by the reaction with oxygen in the presence
of a catalyst (V
2
O
5
), and
(iii) absorption of SO
3
in H
2
SO
4
to give Oleum (H
2
S
2
O
7
).
A flow diagram for the manufacture of sulphuric acid is shown in
(Fig. 7.7). The SO
2
produced is purified by removing dust and other
impurities such as arsenic compounds.
The key step in the manufacture of H
2
SO
4
is the catalytic oxidation
of SO
2
with O
2
to give SO
3
in the presence of V
2
O
5
(catalyst).
(
)
(
)
(
)
2 5
V O
2 2 3
2SO g O g 2SO g
+ 
1
r
196.6 kJmol
= −
0
H
The reaction is exothermic, reversible and the forward reaction leads
to a decrease in volume. Therefore, low temperature and high pressure
are the favourable conditions for maximum yield. But the temperature
should not be very low otherwise rate of reaction will become slow.
In practice, the plant is operated at a pressure of 2 bar and a
temperature of 720 K. The SO
3
gas from the catalytic converter is
absorbed in concentrated H
2
SO
4
to produce oleum. Dilution of oleum
with water gives H
2
SO
4
of the desired concentration. In the industry
two steps are carried out simultaneously to make the process a
continuous one and also to reduce the cost.
SO
3
+ H
2
SO
4
H
2
S
2
O
7
(Oleum)
The sulphuric acid obtained by Contact process is 96-98% pure.
7.177.17
7.177.17
7.17
SulphuricSulphuric
SulphuricSulphuric
Sulphuric
AcidAcid
AcidAcid
Acid
Dust
precipitator
Sulphur
burner
Waste
water
Waste
acid
Drying
tower
Washing and
cooling tower
Arsenic purifier
containing
gelatinous hydrated
ferric oxide
Catalytic
converter
Oleum
(H S O )
2 2 7
SO
3
Conc. H SO
2 4
Dry SO +O
2 2
Water
spray
Conc. H SO
spray
2 4
Impure
SO +O
2 2
Sulphur
Air
Preheater
Quartz
V O
2 5
Fig. 7.7: Flow diagram for the manufacture of sulphuric acid
2020-21
196Chemistry
Properties
Sulphuric acid is a colourless, dense, oily liquid with a specific gravity
of 1.84 at 298 K. The acid freezes at 283 K and boils at 611 K. It
dissolves in water with the evolution of a large quantity of heat. Hence,
care must be taken while preparing sulphuric acid solution from
concentrated sulphuric acid. The concentrated acid must be added
slowly into water with constant stirring.
The chemical reactions of sulphuric acid are as a result of the
following characteristics: (a) low volatility (b) strong acidic character
(c) strong affinity for water and (d) ability to act as an oxidising agent.
In aqueous solution, sulphuric acid ionises in two steps.
H
2
SO
4
(aq) + H
2
O(l) H
3
O
+
(aq) + HSO
4
(aq);
1
a
K
= very large (
1
a
K
>10)
HSO
4
(aq) + H
2
O(l) H
3
O
+
(aq) + SO
4
2-
(aq) ;
2
a
K
= 1.2 × 10
–2
The larger value of
1
a
K
(
1
a
K
>10) means that H
2
SO
4
is largely
dissociated into H
+
and HSO
4
. Greater the value of dissociation constant
(K
a
), the stronger is the acid.
The acid forms two series of salts: normal sulphates (such as sodium
sulphate and copper sulphate) and acid sulphates (e.g., sodium
hydrogen sulphate).
Sulphuric acid, because of its low volatility can be used to
manufacture more volatile acids from their corresponding salts.
2 MX + H
2
SO
4
2 HX + M
2
SO
4
(X = F, Cl, NO
3
)
(M = Metal)
Concentrated sulphuric acid is a strong dehydrating agent. Many
wet gases can be dried by passing them through sulphuric acid,
provided the gases do not react with the acid. Sulphuric acid removes
water from organic compounds; it is evident by its charring action on
carbohydrates.
C
12
H
22
O
11
2 4
H SO

12C + 11H
2
O
Hot concentrated sulphuric acid is a moderately strong oxidising
agent. In this respect, it is intermediate between phosphoric and nitric
acids. Both metals and non-metals are oxidised by concentrated
sulphuric acid, which is reduced to SO
2
.
Cu + 2 H
2
SO
4
(conc.) CuSO
4
+ SO
2
+ 2H
2
O
S + 2H
2
SO
4
(conc.) 3SO
2
+ 2H
2
O
C + 2H
2
SO
4
(conc.) CO
2
+ 2 SO
2
+ 2 H
2
O
UsesUses
UsesUses
Uses: Sulphuric acid is a very important industrial chemical. A nation’s
industrial strength can be judged by the quantity of sulphuric acid it
produces and consumes. It is needed for the manufacture of hundreds
of other compounds and also in many industrial processes. The bulk
of sulphuric acid produced is used in the manufacture of fertilisers
(e.g., ammonium sulphate, superphosphate). Other uses are in:
(a) petroleum refining (b) manufacture of pigments, paints and dyestuff
intermediates (c) detergent industry (d) metallurgical applications (e.g.,
cleansing metals before enameling, electroplating and galvanising
(e) storage batteries (f) in the manufacture of nitrocellulose products
and (g) as a laboratory reagent.
2020-21
197 The p-Block Elements
Fluorine, chlorine, bromine, iodine, astatine and tennessine are
members of Group 17. These are collectively known as the halogens
(Greek halo means salt and genes means born i.e., salt producers).
The halogens are highly reactive non-metallic elements. Like Groups
1 and 2, the elements of Group 17 show great similarity amongst
themselves. That much similarity is not found in the elements of
other groups of the periodic table. Also, there is a regular gradation
in their physical and chemical properties. Astatine and tennessine are
radioactive elements.
Fluorine and chlorine are fairly abundant while bromine and iodine
less so. Fluorine is present mainly as insoluble fluorides (fluorspar
CaF
2
, cryolite Na
3
AlF
6
and fluoroapatite 3Ca
3
(PO
4
)
2
.CaF
2
)
and small
quantities are present in soil, river water plants and bones and teeth
of animals. Sea water contains chlorides, bromides and iodides of
sodium, potassium, magnesium and calcium, but is mainly sodium
chloride solution (2.5% by mass). The deposits of dried up seas
contain these compounds, e.g., sodium chloride and carnallite,
KCl.MgCl
2
.6H
2
O. Certain forms of marine life contain iodine in their
systems; various seaweeds, for example, contain upto 0.5% of iodine
and Chile saltpetre contains upto 0.2% of sodium iodate.
Here important atomic and physical properties of Group 17 elements
other than tennessine are given along with their electronic configurations
[Table 7.8, page 198]. Tennessine is a synthetic radioactive element. Its
symbol is Ts, atomic number 117, atomic mass 294 and electronic
configuration [Rn] 5f
14
6d
10
7s
2
7p
5
. Only very small amount of the element
could be prepared. Also its half life is in milliseconds only. That is why
its chemistry could not be established.
7.18.1 Occurrence
7.187.18
7.187.18
7.18
Group 17Group 17
Group 17Group 17
Group 17
ElementsElements
ElementsElements
Elements
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Intext Questions
7.23 Mention three areas in which H
2
SO
4
plays an important role.
7.24 Write the conditions to maximise the yield of H
2
SO
4
by Contact process.
7.25 Why is
K K
a a
2 1
=
for H
2
SO
4
in water ?
What happens when
(i) Concentrated H
2
SO
4
is added to calcium fluoride
(ii) SO
3
is passed through water?
(i) It forms hydrogen fluoride
2 2 4 4
CaF H SO CaSO 2HF
+ +
(ii) It dissolves SO
3
to give H
2
SO
4
.
3 2 2 4
SO H O H SO
+
SolutionSolution
Solution
Solution
Solution
Example 7.13Example 7.13
Example 7.13Example 7.13
Example 7.13
2020-21
198Chemistry
The trends of some of the atomic, physical and chemical properties
are discussed below.
All these elements have seven electrons in their outermost shell (ns
2
np
5
)
which is one electron short of the next noble gas.
The halogens have the smallest atomic radii in their respective periods
due to maximum effective nuclear charge. The atomic radius of fluorine
like the other elements of second period is extremely small. Atomic and
ionic radii increase from fluorine to iodine due to increasing number
of quantum shells.
They have little tendency to lose electron. Thus they have very high
ionisation enthalpy. Due to increase in atomic size, ionisation enthalpy
decreases down the group.
Halogens have maximum negative electron gain enthalpy in the
corresponding periods. This is due to the fact that the atoms of these
elements have only one electron less than stable noble gas configurations.
Electron gain enthalpy of the elements of the group becomes less negative
down the group. However, the negative electron gain enthalpy of fluorine
is less than that of chlorine. It is due to small size of fluorine atom. As
a result, there are strong interelectronic repulsions in the relatively
small 2p orbitals of fluorine and thus, the incoming electron does not
experience much attraction.
7.18.3 Atomic
and Ionic
Radii
7.18.4 Ionisation
Enthalpy
7.18.5 Electron
Gain
Enthalpy
7.18.2 Electronic
Configuration
Property F Cl Br I At
a
Atomic number 9 17 35 53 85
Atomic mass/g mol
–1
19.00 35.45 79.90 126.90 210
Electronic configuration [He]2s
2
2p
5
[Ne]3s
2
3p
5
[Ar]3d
10
4s
2
4p
5
[Kr]4d
10
5s
2
5p
5
[Xe]4f
14
5d
10
6s
2
6p
5
Covalent radius/pm 64 99 114 133
Ionic radius X
/pm 133 184 196 220
Ionisation enthalpy/kJ mol
–1
1680 1256 1142 1008
Electron gain enthalpy/kJ mol
–1
–333 –349 –325 –296
Electronegativity
b
4 3.2 3.0 2.7 2.2
Hyd
H(X
)/kJ mol
–1
515 381 347 305
F
2
Cl
2
Br
2
I
2
Melting point/K 54.4 172.0 265.8 386.6
Boiling point/K 84.9 239.0 332.5 458.2
Density/g cm
–3
1.5 (85)
c
1.66 (203)
c
3.19(273)
c
4.94(293)
d
Distance X – X/pm 143
199 228 266
Bond dissociation enthalpy 158.8 242.6 192.8 151.1
/(kJ mol
–1
)
E
V
/V
e
2.87
1.36 1.09 0.54
Table 7.8: Atomic and Physical Properties of Halogens
a
Radioactive;
b
Pauling scale;
c
For the liquid at temperatures (K) given in the parentheses;
d
solid;
e
The
half-cell reaction is X
2
(g) + 2e
2X
(aq).
2020-21
199 The p-Block Elements
Oxidation states and trends in chemical reactivity
All the halogens exhibit –1 oxidation state. However, chlorine, bromine
and iodine exhibit + 1, + 3, + 5 and + 7 oxidation states also as
explained below:
7.18.8 Chemical
Properties
Halogen atom
in ground state
(other than fluorine)
1 unpaired electron accounts
for –1 or +1 oxidation states
First excited state
3 unpaired electrons account
for +3 oxidation states
Second excited state
Third excited state
5 unpaired electrons account
for +5 oxidation state
7 unpaired electrons
account for +7 oxidation state
ns np nd
7.18.6
Electronegativity
7.18.7 Physical
Properties
They have very high electronegativity. The electronegativity decreases
down the group. Fluorine is the most electronegative element in the
periodic table.
Halogens have maximum negative electron gain enthalpy in the
respective periods of the periodic table. Why?
Halogens have the smallest size in their respective periods and therefore
high effective nuclear charge. As a consequence, they readily accept
one electron to acquire noble gas electronic configuration.
Example 7.14Example 7.14
Example 7.14Example 7.14
Example 7.14
SolutionSolution
Solution
Solution
Solution
Although electron gain enthalpy of fluorine is less negative as compared
to chlorine, fluorine is a stronger oxidising agent than chlorine. Why?
It is due to
(i) low enthalpy of dissociation of F-F bond (Table 7.8).
(ii) high hydration enthalpy of F
(Table 7.8).
Example 7.15Example 7.15
Example 7.15Example 7.15
Example 7.15
SolutionSolution
SolutionSolution
Solution
Halogens display smooth variations in their physical properties. Fluorine
and chlorine are gases, bromine is a liquid and iodine is a solid. Their
melting and boiling points steadily increase with atomic number. All
halogens are coloured. This is due to absorption of radiations in visible
region which results in the excitation of outer electrons to higher energy
level. By absorbing different quanta of radiation, they display different
colours. For example, F
2
, has yellow, Cl
2
, greenish yellow, Br
2
, red and
I
2
, violet colour. Fluorine and chlorine react with water. Bromine and
iodine are only sparingly soluble in water but are soluble in various
organic solvents such as chloroform, carbon tetrachloride, carbon
disulphide and hydrocarbons to give coloured solutions.
One curious anomaly we notice from Table 7.8 is the smaller
enthalpy of dissociation of F
2
compared to that of Cl
2
whereas X-X
bond dissociation enthalpies from chlorine onwards show the expected
trend: Cl – Cl > Br – Br > I – I. A reason for this anomaly is the relatively
large electron-electron repulsion among the lone pairs in F
2
molecule
where they are much closer to each other than in case of Cl
2
.
2020-21
200Chemistry
The higher oxidation states of chlorine, bromine and iodine are realised
mainly when the halogens are in combination with the small and highly
electronegative fluorine and oxygen atoms. e.g., in interhalogens, oxides
and oxoacids. The oxidation states of +4 and +6 occur in the oxides and
oxoacids of chlorine and bromine. The fluorine atom has no d orbitals
in its valence shell and therefore cannot expand its octet. Being the most
electronegative, it exhibits only –1 oxidation state.
All the halogens are highly reactive. They react with metals and
non-metals to form halides. The reactivity of the halogens decreases
down the group.
The ready acceptance of an electron is the reason for the strong
oxidising nature of halogens. F
2
is the strongest oxidising halogen and
it oxidises other halide ions in solution or even in the solid phase. In
general, a halogen oxidises halide ions of higher atomic number.
F
2
+ 2X
2F
+ X
2
(X = Cl, Br or I)
Cl
2
+ 2X
2Cl
+ X
2
(X = Br or I)
Br
2
+ 2I
2Br
+ I
2
The decreasing oxidising ability of the halogens in aqueous solution
down the group is evident from their standard electrode potentials
(Table 7.8) which are dependent on the parameters indicated below:
( ) ( ) ( )
( )
eg
hyd
diss
1/2
2
1
X g X g X g X
aq
2
H
H
H
 
V
V
V
The relative oxidising power of halogens can further be illustrated
by their reactions with water. Fluorine oxidises water to oxygen whereas
chlorine and bromine react with water to form corresponding hydrohalic
and hypohalous acids. The reaction of iodine with water is non-
spontaneous. In fact, I
can be oxidised by oxygen in acidic medium;
just the reverse of the reaction observed with fluorine.
(
)
( )
(
)
(
)
(
)
( )
( )
( ) ( )
( )
( ) ( )
( )
( )
( )
2 2 2
2 2
2 2 2
2F g 2H O 4H 4F O g
aq aq
l
X g H O HX HOXaq aq
l
where X = Cl or Br
4I 4H O g 2I 2H Oaq aq
s
l
+
+
+ + +
+ +
+ + +
Anomalous behaviour of fluorine
Like other elements of p-block present in second period of the periodic
table, fluorine is anomalous in many properties. For example, ionisation
enthalpy, electronegativity, and electrode potentials are all higher for
fluorine than expected from the trends set by other halogens. Also, ionic
and covalent radii, m.p. and b.p., enthalpy of bond dissociation and electron
gain enthalpy are quite lower than expected. The anomalous behaviour of
fluorine is due to its small size, highest electronegativity, low F-F bond
dissociation enthalpy, and non availability of d orbitals in valence shell.
Most of the reactions of fluorine are exothermic (due to the small
and strong bond formed by it with other elements). It forms only one
oxoacid while other halogens form a number of oxoacids. Hydrogen
fluoride is a liquid (b.p. 293 K) due to strong hydrogen bonding.
Hydrogen bond is formed in HF due to small size and high
2020-21
201 The p-Block Elements
electronegativity of fluorine. Other hydrogen halides which have bigger
size and less electronegativity are gases.
(i) Reactivity towards hydrogen: They all react with hydrogen to give
hydrogen halides but affinity for hydrogen decreases from fluorine
to iodine. Hydrogen halides dissolve in water to form hydrohalic
acids. Some of the properties of hydrogen halides are given in
Table 7.9. The acidic strength of these acids varies in the order:
HF < HCl < HBr < HI. The stability of these halides decreases down
the group due to decrease in bond (H–X) dissociation enthalpy in
the order: H–F > H–Cl > H–Br > H–I.
Property HF HCl HBr HI
Melting point/K 190 159
185 222
Boiling point/K 293 189 206 238
Bond length (H – X)/pm 91.7 127.4 141.4 160.9
diss
H
V
/kJ mol
–1
574 432 363 295
pK
a
3.2 –7.0
–9.5 –10.0
Table 7.9: Properties of Hydrogen Halides
(ii) Reactivity towards oxygen: Halogens form many oxides with oxygen
but most of them are unstable. Fluorine forms two oxides OF
2
and
O
2
F
2
. However, only OF
2
is thermally stable at 298 K. These oxides
are essentially oxygen fluorides because of the higher
electronegativity of fluorine than oxygen. Both are strong fluorinating
agents. O
2
F
2
oxidises plutonium to PuF
6
and the reaction is used
in removing plutonium as PuF
6
from spent nuclear fuel.
Chlorine, bromine and iodine form oxides in which the oxidation
states of these halogens range from +1 to +7. A combination of kinetic
and thermodynamic factors lead to the generally decreasing order of
stability of oxides formed by halogens, I > Cl > Br. Higher stability of
oxides of iodine is due to greater polarisability of bond between
iodine and oxygen. In the case of chlorine, multiple bond formation
between chlorine and oxygen takes place due to availability of
d–orbitals. This leads to increase in stability. Bromine lacks both the
characteristics hence stability of oxides of bromine is least. The higher
oxides of halogens tend to be more stable than the lower ones.
Chlorine oxides, Cl
2
O, ClO
2
, Cl
2
O
6
and Cl
2
O
7
are highly reactive
oxidising agents and tend to explode. ClO
2
is used as a bleaching
agent for paper pulp and textiles and in water treatment.
The bromine oxides, Br
2
O, BrO
2
, BrO
3
are the least stable
halogen oxides (middle row anomally) and exist only at low
temperatures. They are very powerful oxidising agents.
The iodine oxides, I
2
O
4
, I
2
O
5
, I
2
O
7
are insoluble solids and
decompose on heating. I
2
O
5
is a very good oxidising agent and is
used in the estimation of carbon monoxide.
(iii) Reactivity towards metals: Halogens react with metals to form
metal halides. For example, bromine reacts with magnesium to give
magnesium bromide.
2020-21
202Chemistry
(
)
(
)
(
)
2 2
Mg s Br l MgBr s
+
The ionic character of the halides decreases in the order MF >
MCl > MBr > MI where M is a monovalent metal. If a metal exhibits
more than one oxidation state, the halides in higher oxidation
state will be more covalent than the one in lower oxidation state.
For example, SnCl
4
, PbCl
4
, SbCl
5
and UF
6
are more covalent than
SnCl
2
, PbCl
2
, SbCl
3
and UF
4
respectively.
(iv) Reactivity of halogens towards other halogens: Halogens combine
amongst themselves to form a number of compounds known as
interhalogens of the types XX
, XX
3
, XX
5
and XX
7
where X is a
larger size halogen and X
is smaller size halogen.
Fluorine exhibits only 1 oxidation state whereas other halogens
exhibit + 1, + 3, + 5 and + 7 oxidation states also. Explain.
Fluorine is the most electronegative element and cannot exhibit any positive
oxidation state. Other halogens have d orbitals and therefore, can expand
their octets and show + 1, + 3, + 5 and + 7 oxidation states also.
Example 7.16Example 7.16
Example 7.16Example 7.16
Example 7.16
SolutionSolution
SolutionSolution
Solution
Intext QuestionsIntext Questions
Intext QuestionsIntext Questions
Intext Questions
7.26 Considering the parameters such as bond dissociation enthalpy, electron
gain enthalpy and hydration enthalpy, compare the oxidising power of
F
2
and Cl
2
.
7.27 Give two examples to show the anomalous behaviour of fluorine.
7.28 Sea is the greatest source of some halogens. Comment.
Chlorine was discovered in 1774 by Scheele by the action of HCl on
MnO
2
. In 1810 Davy established its elementary nature and suggested the
name chlorine on account of its colour (Greek, chloros = greenish yellow).
Preparation
It can be prepared by any one of the following methods:
(i) By heating manganese dioxide with concentrated hydrochloric acid.
MnO
2
+ 4HCl MnCl
2
+ Cl
2
+ 2H
2
O
However, a mixture of common salt and concentrated H
2
SO
4
is
used in place of HCl.
4NaCl + MnO
2
+ 4H
2
SO
4
MnCl
2
+ 4NaHSO
4
+ 2H
2
O + Cl
2
(ii) By the action of HCl on potassium permanganate.
2KMnO
4
+ 16HCl 2KCl + 2MnCl
2
+ 8H
2
O + 5Cl
2
Manufacture of chlorine
(i) Deacon’s process: By oxidation of hydrogen chloride gas by
atmospheric oxygen in the presence of CuCl
2
(catalyst) at 723 K.
2
CuCl
2 2 2
4HCl O 2Cl 2H O
+  +
(ii) Electrolytic process: Chlorine is obtained by the electrolysis of
brine (concentrated NaCl solution). Chlorine is liberated at anode.
It is also obtained as a by–product in many chemical industries.
7.197.19
7.197.19
7.19
ChlorineChlorine
ChlorineChlorine
Chlorine
2020-21
203 The p-Block Elements
Properties
It is a greenish yellow gas with pungent and suffocating odour. It is
about 2-5 times heavier than air. It can be liquefied easily into greenish
yellow liquid which boils at 239 K. It is soluble in water.
Chlorine reacts with a number of metals and non-metals to form chlorides.
2Al + 3Cl
2
2AlCl
3
; P
4
+ 6Cl
2
4PCl
3
2Na + Cl
2
2NaCl; S
8
+ 4Cl
2
4S
2
Cl
2
2Fe + 3Cl
2
2FeCl
3
;
It has great affinity for hydrogen. It reacts with compounds
containing hydrogen to form HCl.
2 2
2 2
10 16 2
H Cl 2HCl
H S Cl 2HCl S
C H 8Cl 16HCl 10C
+
+ +
+ +
With cold and dilute alkalies chlorine produces a mixture of chloride
and hypochlorite but with hot and concentrated alkalies it gives chloride
and chlorate.
2NaOH + Cl
2
NaCl + NaOCl + H
2
O
(cold and dilute)
6 NaOH + 3Cl
2
5NaCl + NaClO
3
+ 3H
2
O
(hot and conc.)
With dry slaked lime it gives bleaching powder.
2Ca(OH)
2
+ 2Cl
2
Ca(OCl)
2
+ CaCl
2
+ 2H
2
O
The composition of bleaching powder is Ca(OCl)
2
.CaCl
2
.Ca(OH)
2
.2H
2
O.
Chlorine reacts with hydrocarbons and gives substitution products
with saturated hydrocarbons and addition products with unsaturated
hydrocarbons. For example,
CH
4
+ Cl
2
UV
CH
3
Cl + HCl
Methane Methyl chloride
C
2
H
4
+ Cl
2
Room temp.

C
2
H
4
Cl
2
Ethene 1,2-Dichloroethane
Chlorine water on standing loses its yellow colour due to the
formation of HCl and HOCl. Hypochlorous acid (HOCl) so formed, gives
nascent oxygen which is responsible for oxidising and bleaching
properties of chlorine.
Chlorine oxidises ferrous to ferric and sulphite to sulphate. Chlorine
oxidises sulphur dioxide to sulphur trioxide and iodine to iodate. In the
presence of water they form sulphuric acid and iodic acid respectively.
2FeSO
4
+ H
2
SO
4
+ Cl
2
Fe
2
(SO
4
)
3
+ 2HCl
Na
2
SO
3
+ Cl
2
+ H
2
O Na
2
SO
4
+ 2HCl
SO
2
+ 2H
2
O + Cl
2
H
2
SO
4
+ 2HCl
I
2
+ 6H
2
O + 5Cl
2
2HIO
3
+ 10HCl
Chlorine is a powerful bleaching agent; bleaching action is due
to oxidation.
It bleaches vegetable or organic matter in the presence of moisture.
Bleaching effect of chlorine is permanent.
Cl
2
+ H
2
O 2HCl + O
Coloured substance + O Colourless substance
2020-21
204Chemistry
Write the balanced chemical equation for the reaction of Cl
2
with hot
and concentrated NaOH. Is this reaction a disproportionation
reaction? Justify.
3Cl
2
+ 6NaOH 5NaCl + NaClO
3
+ 3H
2
O
Yes, chlorine from zero oxidation state is changed to –1 and +5
oxidation states.
Example 7.17Example 7.17
Example 7.17Example 7.17
Example 7.17
SolutionSolution
Solution
Solution
Solution
Glauber prepared this acid in 1648 by heating common salt with
concentrated sulphuric acid. Davy in 1810 showed that it is a compound
of hydrogen and chlorine.
Preparation
In laboratory, it is prepared by heating sodium chloride with
concentrated sulphuric acid.
NaCl + H
2
SO
4
420 K

NaHSO
4
+ HCl
NaHSO
4
+ NaCl
823K

Na
2
SO
4
+ HCl
HCl gas can be dried by passing through concentrated sulphuric acid.
Properties
It is a colourless and pungent smelling gas. It is easily liquefied to a
colourless liquid (b.p.189 K) and freezes to a white crystalline solid
(f.p. 159 K). It is extremely soluble in water and ionises as follows:
(
)
( )
(
)
(
)
7
2 3 a
HCl g H O H O Cl 10
aq aq
l
+
+ + =K
Its aqueous solution is called hydrochloric acid. High value of
dissociation constant (K
a
) indicates that it is a strong acid in water.
It reacts with NH
3
and gives white fumes of NH
4
Cl.
NH
3
+ HCl NH
4
Cl
When three parts of concentrated HCl and one part of concentrated
HNO
3
are mixed, aqua regia is formed which is used for dissolving
noble metals, e.g., gold, platinum.
3 4 2
2
3 6 2
Au 4H NO 4Cl AuCl NO 2H O
3Pt 16H 4NO 18Cl 3PtCl 4NO 8H O
+
+
+ + + + +
+ + + + +
7.207.20
7.207.20
7.20
HydrogenHydrogen
HydrogenHydrogen
Hydrogen
ChlorideChloride
ChlorideChloride
Chloride
UsesUses
UsesUses
Uses: It is used (i) for bleaching woodpulp (required for the manufacture of
paper and rayon), bleaching cotton and textiles, (ii) in the extraction of gold
and platinum (iii) in the manufacture of dyes, drugs and organic compounds
such as CCl
4
, CHCl
3
, DDT, refrigerants, etc. (iv) in sterilising drinking water
and (v) preparation of poisonous gases such as phosgene (COCl
2
), tear gas
(CCl
3
NO
2
), mustard gas (ClCH
2
CH
2
SCH
2
CH
2
Cl).
Intext QuestionsIntext Questions
Intext QuestionsIntext Questions
Intext Questions
7.29 Give the reason for bleaching action of Cl
2
.
7.30 Name two poisonous gases which can be prepared from chlorine gas.
2020-21
205 The p-Block Elements
Fig. 7.8
The structures of
oxoacids of chlorine
Hydrochloric acid decomposes salts of weaker acids, e.g.,
carbonates, hydrogencarbonates, sulphites, etc.
Na
2
CO
3
+ 2HCl 2NaCl + H
2
O + CO
2
NaHCO
3
+ HCl NaCl + H
2
O + CO
2
Na
2
SO
3
+ 2HCl 2NaCl + H
2
O + SO
2
Table 7.10: Oxoacids of Halogens
Halic (I) acid HOF HOCl HOBr HOI
(Hypohalous acid) (Hypofluorous acid) (Hypochlorous acid) (Hypobromous acid) (Hypoiodous acid)
Halic (III) acid HOCIO
(Halous acid) (chlorous acid)
Halic (V) acid HOCIO
2
HOBrO
2
HOIO
2
(Halic acid) (chloric acid) (bromic acid) (iodic acid)
Halic (VII) acid HOCIO
3
HOBrO
3
HOIO
3
(Perhalic acid) (perchloric acid) (perbromic acid) (periodic acid)
When HCl reacts with finely powdered iron, it forms ferrous chloride
and not ferric chloride. Why?
Its reaction with iron produces H
2
.
2 2
Fe 2HCl FeCl H
+ +
Liberation of hydrogen prevents the formation of ferric chloride.
Example 7.18Example 7.18
Example 7.18Example 7.18
Example 7.18
SolutionSolution
SolutionSolution
Solution
UsesUses
UsesUses
Uses: It is used (i) in the manufacture of chlorine, NH
4
Cl and glucose (from
corn starch), (ii) for extracting glue from bones and purifying bone black, (iii)
in medicine and as a laboratory reagent.
Due to high electronegativity and small size, fluorine forms only one
oxoacid, HOF known as fluoric (I) acid or hypofluorous acid. The other
halogens form several oxoacids. Most of them cannot be isolated in
pure state. They are stable only in aqueous solutions or in the form of
their salts. The oxoacids of halogens are given in Table 7.10 and their
structures are given in Fig. 7.8.
7.21 Oxoacids of7.21 Oxoacids of
7.21 Oxoacids of7.21 Oxoacids of
7.21 Oxoacids of
HalogensHalogens
HalogensHalogens
Halogens
2020-21
206Chemistry
When two different halogens react with each other, interhalogen
compounds are formed. They can be assigned general compositions as
XX
, XX
3
, XX
5
and XX
7
where X is halogen of larger size and X
of
smaller size and X is more electropositive than X
. As the ratio between
radii of X and X
increases, the number of atoms per molecule also
increases. Thus, iodine (VII) fluoride should have maximum number of
atoms as the ratio of radii between I and F should be maximum. That
is why its formula is IF
7
(having maximum number of atoms).
Preparation
The interhalogen compounds can be prepared by the direct
combination or by the action of halogen on lower interhalogen
compounds. The product formed depends upon some specific
conditions, For example,
7.227.22
7.227.22
7.22
InterhalogenInterhalogen
InterhalogenInterhalogen
Interhalogen
CompoundsCompounds
CompoundsCompounds
Compounds
2 2
(equimolar)
I Cl 2ICl;
+
573 K
2 2 3
(excess)
Cl 3F 2ClF ;
+ 
437 K
2 2
(equal volume)
Cl F 2ClF;
+ 
2 2 3
(excess)
I 3Cl 2ICl
+
2 2 3
(diluted with water)
Br 3F 2BrF
+
2 2 5
(excess)
Br 5F 2BrF
+
Properties
Some properties of interhalogen compounds are given in Table 7.11.
Type Formula Physical state and colour Structure
XX
1
ClF colourless gas
BrF pale brown gas
IF
a
detected spectroscopically
BrCl
b
gas
ICl ruby red solid (α-form)
brown red solid (β-form)
IBr black solid
XX
3
ClF
3
colourless gas Bent T-shaped
BrF
3
yellow green liquid Bent T-shaped
IF
3
yellow powder Bent T-shaped (?)
ICl
3
c
orange solid Bent T-shaped (?)
XX
5
IF
5
colourless gas but Square
solid below 77 K pyramidal
BrF
5
colourless liquid Square
pyramidal
ClF
5
colourless liquid Square
pyramidal
XX
7
IF
7
colourless gas Pentagonal
bipyramidal
Table 7.11: Some Properties of Interhalogen Compounds
a
Very unstable;
b
The pure solid is known at room temperature;
c
Dimerises as Cl–bridged
dimer (I
2
Cl
6
)
2020-21
207 The p-Block Elements
These are all covalent molecules and are diamagnetic in nature.
They are volatile solids or liquids at 298 K except ClF which is a
gas. Their physical properties are intermediate between those of
constituent halogens except that their m.p. and b.p. are a little higher
than expected.
Their chemical reactions can be compared with the individual
halogens. In general, interhalogen compounds are more reactive
than halogens (except fluorine). This is because X–X bond in
interhalogens is weaker than X–X bond in halogens except F–F
bond. All these undergo hydrolysis giving halide ion derived from
the smaller halogen and a hypohalite ( when XX), halite ( when
XX
3
), halate (when XX
5
) and perhalate (when XX
7
) anion derived
from the larger halogen.
2
XX H O HX HOX
' '
+ +
Their molecular structures are very interesting which can be
explained on the basis of VSEPR theory (Example 7.19). The XX
3
compounds have the bent ‘T’ shape, XX
5
compounds square pyramidal
and IF
7
has pentagonal bipyramidal structures (Table 7.11).
Discuss the molecular shape of BrF
3
on the basis of VSEPR theory.
The central atom Br has seven electrons
in the valence shell. Three of these will form electron-
pair bonds with three fluorine atoms leaving behind
four electrons. Thus, there are three bond pairs and
two lone pairs. According to VSEPR theory, these
will occupy the corners of a trigonal bipyramid. The
two lone pairs will occupy the equatorial positions
to minimise lone pair-lone pair and the bond pair-
lone pair repulsions which are greater than the bond
pair-bond pair repulsions. In addition, the axial
fluorine atoms will be bent towards the equatorial
fluorine in order to minimise the lone-pair-lone pair
repulsions. The shape would be that of a slightly
bent ‘T’.
Example 7.19Example 7.19
Example 7.19Example 7.19
Example 7.19
SolutionSolution
SolutionSolution
Solution
UsesUses
UsesUses
Uses: These compounds can be used as non aqueous solvents. Interhalogen
compounds are very useful fluorinating agents. ClF
3
and BrF
3
are used for the
production of UF
6
in the enrichment of
235
U.
U(s) + 3ClF
3
(l) UF
6
(g) + 3ClF(g)
Intext QuestionIntext Question
Intext QuestionIntext Question
Intext Question
7.31 Why is ICl more reactive than I
2
?
2020-21
208Chemistry
Group 18 consists of elements: helium, neon, argon, krypton, xenon,
radon and oganesson. All these are gases and chemically unreactive.
They form very few compounds, because of this they are termed
as noble gases.
All these gases except radon and oganesson occur in the atmosphere.
Their atmospheric abundance in dry air is ~ 1% by volume of which
argon is the major constituent. Helium and sometimes neon are found
in minerals of radioactive origin e.g., pitchblende, monazite, cleveite.
The main commercial source of helium is natural gas. Xenon and
radon are the rarest elements of the group. Radon is obtained as a
decay product of
226
Ra.
226 222 4
88 86 2
Ra Rn He
+
Oganesson has been synthetically produced by collision of
249
98
Cf
atoms and
48
20
Ca
ions
249 48 294
98 20 118
Cf Ca Og + 3
+ n
7.237.23
7.237.23
7.23
Group 18Group 18
Group 18Group 18
Group 18
ElementsElements
ElementsElements
Elements
7.23.1 Occurrence
Why are the elements of Group 18 known as noble gases ?
The elements present in Group 18 have their valence shell orbitals
completely filled and, therefore, react with a few elements only under
certain conditions. Therefore, they are now known as noble gases.
Example 7.20Example 7.20
Example 7.20Example 7.20
Example 7.20
SolutionSolution
SolutionSolution
Solution
Oganesson has its symbol Og, atomic number 118, atomic mass
294 and electronic configuration [Rn] 5f
14
6d
10
7s
2
7p
6
. Only very small
amount of Og has been produced. Its half life is 0.7 milliseconds.
Therefore, mainly predictions about its chemistry have been made.
Here, except for oganesson, important atomic and physical
properties of other elements of Group 18 along with their electronic
configurations are given in Table 7.12. The trends in some of the atomic,
physical and chemical properties of the group are discussed here.
Propery He Ne Ar Kr Xe Rn*
Atomic number 2 10 18 36 54 86
Atomic mass/ g mol
–1
4.00 20.18 39.95 83.80 131.30 222.00
Electronic configuration 1s
2
[He]2s
2
2p
6
[Ne] 3s
2
3p
6
[Ar]3d
10
4s
2
4p
6
[Kr]4d
10
5s
2
5p
6
[Xe]4f
14
5d
10
6s
2
6p
6
Atomic radius/pm 120 160 190 200 220
Ionisation enthalpy 2372 2080 1520 1351 1170 1037
/kJmol
-1
Electron gain enthalpy 48 116 96 96 77 68
/kJmol
-1
Density (at STP)/gcm
–3
1.8×10
–4
9.0×10
–4
1.8×10
–3
3.7×10
–3
5.9×10
–3
9.7×10
–3
Melting point/K 24.6
83.8 115.9 161.3 202
Boiling point/K 4.2 27.1 87.2 119.7 165.0 211
Atmospheric content 5.24×10
–4
1.82×10
–3
0.934 1.14×10
–4
8.7×10
–6
(% by volume)
Table 7.12: Atomic and Physical Properties of Group 18 Elements
* radioactive
2020-21
209 The p-Block Elements
All noble gases have general electronic configuration ns
2
np
6
except
helium which has 1s
2
(Table 7.12). Many of the properties of noble
gases including their inactive nature are ascribed to their closed
shell structures.
Due to stable electronic configuration these gases exhibit very high
ionisation enthalpy. However, it decreases down the group with increase
in atomic size.
Atomic radii increase down the group with increase in atomic
number.
Since noble gases have stable electronic configurations, they have no
tendency to accept the electron and therefore, have large positive values
of electron gain enthalpy.
Physical Properties
All the noble gases are monoatomic. They are colourless, odourless
and tasteless. They are sparingly soluble in water. They have very low
melting and boiling points because the only type of interatomic
interaction in these elements is weak dispersion forces. Helium has the
lowest boiling point (4.2 K) of any known substance. It has an unusual
property of diffusing through most commonly used laboratory materials
such as rubber, glass or plastics.
7.23.2 Electronic
Configuration
7.23.3 Ionisation
Enthalpy
7.23.4 Atomic
Radii
7.23.5 Electron
Gain
Enthalpy
Chemical Properties
In general, noble gases are least reactive. Their inertness to chemical
reactivity is attributed to the following reasons:
(i) The noble gases except helium (1s
2
) have completely filled ns
2
np
6
electronic configuration in their valence shell.
(ii) They have high ionisation enthalpy and more positive electron
gain enthalpy.
The reactivity of noble gases has been investigated occasionally,
ever since their discovery, but all attempts to force them to react to
form the compounds, were unsuccessful for quite a few years. In March
1962, Neil Bartlett, then at the University of British Columbia, observed
the reaction of a noble gas. First, he prepared a red compound which
is formulated as O
2
+
PtF
6
. He, then realised that the first ionisation
enthalpy of molecular oxygen (1175 kJmol
–1
) was almost identical with
that of xenon (1170 kJ mol
–1
). He made efforts to prepare same type of
compound with Xe and was successful in preparing another red colour
compound Xe
+
PtF
6
by mixing PtF
6
and xenon. After this discovery, a
number of xenon compounds mainly with most electronegative elements
like fluorine and oxygen, have been synthesised.
The compounds of krypton are fewer. Only the difluoride (KrF
2
) has
been studied in detail. Compounds of radon have not been isolated
Noble gases have very low boiling points. Why?
Noble gases being monoatomic have no interatomic forces except weak
dispersion forces and therefore, they are liquefied at very low
temperatures. Hence, they have low boiling points.
Example 7.21Example 7.21
Example 7.21Example 7.21
Example 7.21
SolutionSolution
SolutionSolution
Solution
2020-21
210Chemistry
Fig. 7.9
The structures of
(a) XeF
2
(b) XeF
4
(c) XeF
6
(d) XeOF
4
and (e) XeO
3
F
F
Xe
(a) Linear
(b) Square planar
F
F
F
F
Xe
Xe
F
O
F
F
F
O
O
O
Xe
Xe
F
F
F
F
F
F
(c) Distorted octahedral (d) Square pyramidal
(e) Pyramidal
but only identified (e.g., RnF
2
) by radiotracer technique. No true
compounds of Ar, Ne or He are yet known.
(a) Xenon-fluorine compounds
Xenon forms three binary fluorides, XeF
2
, XeF
4
and XeF
6
by the
direct reaction of elements under appropriate experimental conditions.
Xe (g) + F
2
(g)
673 K, 1 bar

XeF
2
(s)
(xenon in excess)
Xe (g) + 2F
2
(g)
873 K, 7 bar

XeF
4
(s)
(1:5 ratio)
Xe (g) + 3F
2
(g)
573 K , 60 70bar

XeF
6
(s)
(1:20 ratio)
XeF
6
can also be prepared by the interaction of XeF
4
and O
2
F
2
at 143K.
4 2 2 6 2
XeF O F XeF O
+ +
XeF
2
, XeF
4
and XeF
6
are colourless crystalline solids and sublime
readily at 298 K. They are powerful fluorinating agents. They are readily
hydrolysed even by traces of water. For example, XeF
2
is hydrolysed to
give Xe, HF and O
2
.
2XeF
2
(s) + 2H
2
O(l) 2Xe (g) + 4 HF(aq) + O
2
(g)
The structures of the three xenon fluorides can be deduced from
VSEPR and these are shown in Fig. 7.9. XeF
2
and XeF
4
have linear and
square planar structures respectively. XeF
6
has seven electron pairs (6
bonding pairs and one lone pair) and would, thus, have a distorted
octahedral structure as found experimentally in the gas phase.
Xenon fluorides react with fluoride ion acceptors to form cationic
species and fluoride ion donors to form fluoroanions.
XeF
2
+ PF
5
[XeF]
+
[PF
6
]
; XeF
4
+ SbF
5
[XeF
3
]
+
[SbF
6
]
XeF
6
+ MF M
+
[XeF
7
]
(M = Na, K, Rb or Cs)
(b) Xenon-oxygen compounds
Hydrolysis of XeF
4
and XeF
6
with water gives Xe0
3
.
6XeF
4
+ 12 H
2
O 4Xe + 2Xe0
3
+ 24 HF + 3 O
2
XeF
6
+ 3 H
2
O XeO
3
+ 6 HF
Partial hydrolysis of XeF
6
gives oxyfluorides, XeOF
4
and XeO
2
F
2
.
XeF
6
+ H
2
O XeOF
4
+ 2 HF
XeF
6
+ 2 H
2
O XeO
2
F
2
+ 4HF
2020-21
211 The p-Block Elements
XeO
3
is a colourless explosive solid and has a pyramidal molecular
structure (Fig. 7.9). XeOF
4
is a colourless volatile liquid and has a
square pyramidal molecular structure (Fig.7.9).
Does the hydrolysis of XeF
6
lead to a redox reaction?
No, the products of hydrolysis are XeOF
4
and XeO
2
F
2
where the oxidation
states of all the elements remain the same as it was in the reacting state.
Example 7.22Example 7.22
Example 7.22Example 7.22
Example 7.22
SolutionSolution
Solution
Solution
Solution
Uses
Uses
Uses
Uses
Uses: Helium is a non-inflammable and light gas. Hence, it is used in filling
balloons for meteorological observations. It is also used in gas-cooled nuclear
reactors. Liquid helium (b.p. 4.2 K) finds use as cryogenic agent for carrying out
various experiments at low temperatures. It is used to produce and sustain
powerful superconducting magnets which form an essential part of modern NMR
spectrometers and Magnetic Resonance Imaging (MRI) systems for clinical
diagnosis. It is used as a diluent for oxygen in modern diving apparatus because
of its very low solubility in blood.
Neon is used in discharge tubes and fluorescent bulbs for advertisement display
purposes. Neon bulbs are used in botanical gardens and in green houses.
Argon is used mainly to provide an inert atmosphere in high temperature
metallurgical processes (arc welding of metals or alloys) and for filling electric bulbs.
It is also used in the laboratory for handling substances that are air-sensitive.
There are no significant uses of Xenon and Krypton. They are used in light
bulbs designed for special purposes.
Intext QuestionsIntext Questions
Intext QuestionsIntext Questions
Intext Questions
7.32 Why is helium used in diving apparatus?
7.33 Balance the following equation: XeF
6
+ H
2
O XeO
2
F
2
+ HF
7.34 Why has it been difficult to study the chemistry of radon?
2020-21
212Chemistry
SummarySummary
SummarySummary
Summary
Groups 13 to 18 of the periodic table consist of p-block elements with their valence
shell electronic configuration ns
2
np
1–6
. Groups 13 and 14 were dealt with in Class
XI. In this Unit remaining groups of the p-block have been discussed.
Group 15 consists of five elements namely, N, P, As, Sb and Bi which have
general electronic configuration ns
2
np
3
. Nitrogen differs from other elements of this
group due to small size, formation of p
ππ
ππ
π
p
ππ
ππ
π
multiple bonds with itself and with
highly electronegative atom like O or C and non-availability of d orbitals to expand
its valence shell. Elements of group 15 show gradation in properties. They react with
oxygen, hydrogen and halogens. They exhibit two important oxidation states, + 3
and + 5 but +3 oxidation is favoured by heavier elements due to ‘inert pair effect’.
Dinitrogen can be prepared in laboratory as well as on industrial scale. It forms
oxides in various oxidation states as N
2
O, NO, N
2
O
3
, NO
2
, N
2
O
4
and N
2
O
5
. These
oxides have resonating structures and have multiple bonds. Ammonia can be
prepared on large scale by Haber’s process. HNO
3
is an important industrial
chemical. It is a strong monobasic acid and is a powerful oxidising agent. Metals
and non-metals react with HNO
3
under different conditions to give NO or NO
2
.
Phosphorus exists as P
4
in elemental form. It exists in several allotropic forms.
It forms hydride, PH
3
which is a highly poisonous gas. It forms two types of halides as
PX
3
and PX
5
. PCl
3
is prepared by the reaction of white phosphorus with dry chlorine
while PCl
5
is prepared by the reaction of phosphorus with SO
2
Cl
2
. Phosphorus forms
a number of oxoacids. Depending upon the number of P–OH groups, their basicity
varies. The oxoacids which have P–H bonds are good reducing agents.
The Group 16 elements have general electronic configuration ns
2
np
4
. They show
maximum oxidation state, +6. Gradation in physical and chemical properties is
observed in the group 16 elements. In laboratory, dioxygen is prepared by heating
KClO
3
in presence of MnO
2
. It forms a number of oxides with metals. Allotropic form
of oxygen is O
3
which is a highly oxidising agent. Sulphur forms a number of allotropes.
Of these,
α
– and
β
– forms of sulphur are the most important. Sulphur combines with
oxygen to give oxides such as SO
2
and SO
3
. SO
2
is prepared by the direct union of
sulphur with oxygen. SO
2
is used in the manufacture of H
2
SO
4
. Sulphur forms a
number of oxoacids. Amongst them, the most important is H
2
SO
4
. It is prepared by
contact process. It is a dehydrating and oxidising agent. It is used in the manufacture
of several compounds.
Group 17 of the periodic table consists of the following elements F, Cl, Br, I and
At.These elements are extremely reactive and as such they are found in the
combined state only. The common oxidation state of these elements is –1. However,
highest oxidation state can be +7. They show regular gradation in physical and
chemical properties. They form oxides, hydrogen halides, interhalogen compounds
and oxoacids. Chlorine is conveniently obtained by the reaction of HCl with KMnO
4
.
HCl is prepared by heating NaCl with concentrated H
2
SO
4
. Halogens combine with
one another to form interhalogen compounds of the type XX
1
n
(n = 1, 3, 5, 7)
where X
1
is lighter than X. A number of oxoacids of halogens are known. In the
structures of these oxoacids, halogen is the central atom which is bonded in each
case with one OH bond as X–OH. In some cases X = 0 bonds are also found.
Group 18 of the periodic table consists of noble gases. They have ns
2
np
6
valence
shell electronic configuration except He which has 1s
2
. All the gases except Rn
occur in atmosphere. Rn is obtained as the decay product of
226
Ra.
Due to complete octet of outermost shell, they have less tendency to form
compounds. The best characterised compounds are those of xenon with fluorine
and oxygen only under certain conditions. These gases have several uses. Argon is
used to provide inert atmosphere, helium is used in filling balloons for meteorological
observations, neon is used in discharge tubes and fluorescent bulbs.
2020-21
213 The p-Block Elements
7.1 Discuss the general characteristics of Group 15 elements with reference to
their electronic configuration, oxidation state, atomic size, ionisation enthalpy
and electronegativity.
7.2 Why does the reactivity of nitrogen differ from phosphorus?
7.3 Discuss the trends in chemical reactivity of group 15 elements.
7.4 Why does NH
3
form hydrogen bond but PH
3
does not?
7.5 How is nitrogen prepared in the laboratory? Write the chemical equations
of the reactions involved.
7.6 How is ammonia manufactured industrially?
7.7 Illustrate how copper metal can give different products on reaction with HNO
3
.
7.8 Give the resonating structures of NO
2
and N
2
O
5
.
7.9 The HNH angle value is higher than HPH, HAsH and HSbH angles. Why?
[Hint: Can be explained on the basis of sp
3
hybridisation in NH
3
and only
s–p bonding between hydrogen and other elements of the group].
7.10 Why does R
3
P = O exist but R
3
N = O does not (R = alkyl group)?
7.11 Explain why NH
3
is basic while BiH
3
is only feebly basic.
7.12 Nitrogen exists as diatomic molecule and phosphorus as P
4
. Why?
7.13 Write main differences between the properties of white phosphorus and red
phosphorus.
7.14 Why does nitrogen show catenation properties less than phosphorus?
7.15 Give the disproportionation reaction of H
3
PO
3
.
7.16 Can PCl
5
act as an oxidising as well as a reducing agent? Justify.
7.17 Justify the placement of O, S, Se, Te and Po in the same group of the
periodic table in terms of electronic configuration, oxidation state and hydride
formation.
7.18 Why is dioxygen a gas but sulphur a solid?
7.19 Knowing the electron gain enthalpy values for O O
and O O
2–
as –141
and 702 kJ mol
–1
respectively, how can you account for the formation of a
large number of oxides having O
2–
species and not O
?
(Hint: Consider lattice energy factor in the formation of compounds).
7.20 Which aerosols deplete ozone?
7.21 Describe the manufacture of H
2
SO
4
by contact process?
7.22 How is SO
2
an air pollutant?
7.23 Why are halogens strong oxidising agents?
7.24 Explain why fluorine forms only one oxoacid, HOF.
7.25 Explain why inspite of nearly the same electronegativity, nitrogen forms
hydrogen bonding while chlorine does not.
7.26 Write two uses of ClO
2
.
7.27 Why are halogens coloured?
7.28 Write the reactions of F
2
and Cl
2
with water.
7.29 How can you prepare Cl
2
from HCl and HCl from Cl
2
? Write reactions only.
7.30 What inspired N. Bartlett for carrying out reaction between Xe and PtF
6
?
7.31 What are the oxidation states of phosphorus in the following:
(i) H
3
PO
3
(ii) PCl
3
(iii) Ca
3
P
2
(iv) Na
3
PO
4
(v) POF
3
?
Exercises
2020-21
214Chemistry
7.32 Write balanced equations for the following:
(i) NaCl is heated with sulphuric acid in the presence of MnO
2
.
(ii) Chlorine gas is passed into a solution of NaI in water.
7.33 How are xenon fluorides XeF
2
, XeF
4
and XeF
6
obtained?
7.34 With what neutral molecule is ClO
isoelectronic? Is that molecule a Lewis
base?
7.35 How are XeO
3
and XeOF
4
prepared?
7.36 Arrange the following in the order of property indicated for each set:
(i) F
2
, Cl
2
, Br
2
, I
2
- increasing bond dissociation enthalpy.
(ii) HF, HCl, HBr, HI - increasing acid strength.
(iii) NH
3
, PH
3
, AsH
3
, SbH
3
, BiH
3
– increasing base strength.
7.37 Which one of the following does not exist?
(i) XeOF
4
(ii) NeF
2
(iii) XeF
2
(iv) XeF
6
7.38 Give the formula and describe the structure of a noble gas species which
is isostructural with:
(i) ICl
4
(ii) IBr
2
(iii) BrO
3
7.39 Why do noble gases have comparatively large atomic sizes?
7.40 List the uses of neon and argon gases.
Answers to Some Intext Questions
7.1 Higher the positive oxidation state of central atom, more will be its polarising
power which, in turn, increases the covalent character of bond formed
between the central atom and the other atom.
7.2 Because BiH
3
is the least stable among the hydrides of Group 15.
7.3 Because of strong pπ–pπ overlap resulting into the triple bond, NN.
7.6 From the structure of N
2
O
5
it is evident that covalence of nitrogen is four.
7.7(a) Both are sp
3
hybridised. In PH
4
+
all the four orbitals are bonded whereas
in PH
3
there is a lone pair of electrons on P, which is responsible for lone
pair-bond pair repulsion in PH
3
reducing the bond angle to less than
109° 28.
7.10 PCl
5
+ H
2
O POCl
3
+ 2HCl
7.11 Three P–OH groups are present in the molecule of H
3
PO
4
. Therefore, its
basicity is three.
7.15 Because of small size and high electronegativity of oxygen, molecules of
water are highly associated through hydrogen bonding resulting in its
liquid state.
7.21 Both the S–O bonds are covalent and have equal strength due to resonating
structures.
7.25 H
2
SO
4
is a very strong acid in water largely because of its first ionisation
to H
3
O
+
and HSO
4
. The ionisation of HSO
4
to H
3
O
+
and SO
4
2–
is very very
small. That is why K
a
2
<< K
a
1
.
7.31 In general, interhalogen compounds are more reactive than halogens due
to weaker X–X
1
bonding than X–X bond. Thus, ICl is more reactive than I
2
.
7.34 Radon is radioactive with very short half-life which makes the study of
chemistry of radon difficult.
2020-21