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

1–6

(except He which has 1s

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

-Block

Elements

The

-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

(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

7.1

Group 15Group 15

Group 15

ElementsElements

Elements

7.1.1 Occurrence

UnitUnit

Unit

2020-21

171 The p-Block Elements

of the apatite family, Ca

(PO

)

. CaX

(X = F, Cl or OH) (e.g., fluorapatite

(PO

)

. CaF

) 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

. 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

III

single bond (E = element);

3–

;

White phosphorus;

Grey

-form at 38.6 atm;

Sublimation temperature;

At 63 K;

Grey

-form; * Molecular N

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

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 ∆

< ∆

(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

[Ne]3s

[Ar]3d

[Kr]4d

[Xe]4f

Ionisation enthalpy I 1402 1012 947 834 703

(∆

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

70 110 121 141 148

Ionic radius/pm 171

212

222

103

Melting point/K 63* 317

1089

904 544

Boiling point/K 77.2* 554

888

1860 1837

Density/[g cm

–3

(298 K)] 0.879

1.823 5.778

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

. 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

→ HNO

+ H

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

–

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

ππ

-p

ππ

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

-p

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

–p

ππ

bond as the heavier

elements can e.g., R

P = O or R

P = CH

(R = alkyl group). Phosphorus

and arsenic can form d

ππ

–d

ππ

bond also with transition metals when

their compounds like P(C

)

and As(C

)

act as ligands.

(i) Reactivity towards hydrogen: All the elements of Group 15

form hydrides of the type EH

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

to BiH

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

is the

strongest reducing agent amongst all the hydrides. Basicity also

decreases in the order NH

> PH

> AsH

> SbH

> BiH

. Due to

high electronegativity and small size of nitrogen, NH

exhibits

hydrogen bonding in solid as well as liquid state. Because of this,

it has higher melting and boiling points than that of PH

(ii) Reactivity towards oxygen: All these elements form two types

of oxides: E

and E

. 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

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

and EX

. 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

AsH

SbH

BiH

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 –

∆

/kJ mol

–1

–46.1 13.4

66.4 145.1 278

∆

diss

(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.

has lower boiling point than NH

. Why?

Unlike NH

, PH

molecules are not associated through hydrogen bonding

in liquid state. That is why the boiling point of PH

is lower than NH

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.

CI(aq) + NaNO

(aq)

→

(g) + 2H

O(l) + NaCl (aq)

Small amounts of NO and HNO

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

)

Heat

→

+ 4H

O + Cr

Very pure nitrogen can be obtained by the thermal decomposition

of sodium or barium azide.

Ba(N

)

→

Ba + 3N

7.27.2

7.2

DinitrogenDinitrogen

Dinitrogen

Example 7.1Example 7.1

Example 7.1

SolutionSolution

Solution

Example 7.2Example 7.2

Example 7.2

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

is known to be stable. Trihalides

except BiF

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

(calcium nitride) Ca

(calcium phosphide),

As (sodium arsenide), Zn

(zinc antimonide) and Mg

(magnesium bismuthide).

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

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.

NH CONH H O NH CO NH H O CO

2 2 2 4

3 3 2 2

2 2+ →

(

)

+ +

On a small scale ammonia is obtained from ammonium salts which

decompose when treated with caustic soda or calcium hydroxide.

2NH

Cl + Ca(OH)

→ 2NH

+ 2H

O + CaCl

(NH

)

+ 2NaOH → 2NH

+ 2H

O + Na

Properties

Dinitrogen is a colourless, odourless, tasteless and non-toxic gas.

Nitrogen atom has two stable isotopes:

N and

N. It has a very low

solubility in water (23.2 cm

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 N≡N 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

Heat

→

2Li

3Mg + N

Heat

→

It combines with hydrogen at about 773 K in the presence of a

catalyst (Haber’s Process) to form ammonia:

(g) + 3H

(g)

773 k

 

 

2NH

(g); ∆

= –46.1 kJmol

–1

Dinitrogen combines with dioxygen only at very high temperature

(at about 2000 K) to form nitric oxide, NO.

+ O

(g)

Heat

 

 

2NO(g)

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 7.3

SolutionSolution

Solution

Intext QuestionIntext Question

Intext Question

7.3 Why is N

less reactive at room temperature?

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.

(g) + H

O(l) l NH

(aq) + OH

–

(aq)

It forms ammonium salts with acids, e.g., NH

Cl, (NH

)

, 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

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

brown ppt

+ → +x

Fig. 7.1

Flow chart for the

manufacture of

ammonia

On a large scale, ammonia is manufactured by Haber’s process.

(g) + 3H

(g) Ö 2NH

(g); ∆

= – 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

Pa (about 200

atm), a temperature of ~ 700 K and the use of a catalyst such as iron

oxide with small amounts of K

O and Al

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 7.4

Why does NH

act as a Lewis base ?

Nitrogen atom in NH

has one lone pair of electrons which

is available for donation. Therefore, it acts as a Lewis base.

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

7.47.4

7.4

Oxides ofOxides of

Oxides of

NitrogenNitrogen

Nitrogen

SolutionSolution

Solution

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

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

Fe SO 2NaHSO

→ +

neutral

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

, Ag

(aq) + 4 NH

(aq) Ö [Cu(NH

)

]

(aq)

(blue) (deep blue)

(

)

(

)

( )

Ag Cl AgCl

aq aq

+ −

+ →

(colourless) (white ppt)

( )

(

)

(

)

(

)

NHAg

AgCl 2NH Cl

aq aq

 + →

 

(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

+ 3

250 K

2 4 2 3

2NO N O 2N O

+ →

blue solid,

[Nitrogen(III) oxide] acidic

Nitrogen dioxide NO

+ 4

(

)

→

+ +

673K

2 2

2Pb NO

4NO 2PbO O

brown gas,

[Nitrogen(IV) oxide] acidic

Dinitrogen tetroxide N

+ 4

Cool

2 2 4

Heat

2NO N O





colourless solid/

liquid, acidic

[Nitrogen(IV) oxide]

Dinitrogen pentoxide N

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

or NaNO

and concentrated H

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

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

2NO g O g 2NO g



Nitrogen dioxide so formed, dissolves in water to give HNO

(

)

(

)

(

)

(

)

2 2 3

3NO g H O l 2HNO aq NO g

+ → +

NO thus formed is recycled and the aqueous HNO

can be

concentrated by distillation upto ~ 68% by mass. Further

concentration to 98% can be achieved by dehydration with

concentrated H

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

by mass and has a

specific gravity of 1.504.

In the gaseous state, HNO

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

(aq) + H

O(l) → H

(aq) + NO

–

(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

dimerise ?

contains odd number of valence electrons. It behaves as a typical

odd molecule. On dimerisation, it is converted to stable N

molecule

with even number of electrons.

Intext QuestionIntext Question

Intext Question

7.6 What is the covalence of nitrogen in N

7.5

Nitric Acid

Nitrogen forms oxoacids such as H

(hyponitrous acid), HNO

(nitrous acid) and HNO

(nitric acid). Amongst them HNO

is the

most important.

SolutionSolution

Solution

Example 7.5Example 7.5

Example 7.5

2020-21

180Chemistry

7.67.6

7.6

Phosphorus —Phosphorus —

Phosphorus —

AllotropicAllotropic

Allotropic

FormsForms

Forms

products of oxidation depend upon the concentration of the acid,

temperature and the nature of the material undergoing oxidation.

3Cu + 8 HNO

(dilute) → 3Cu(NO

)

+ 2NO + 4H

Cu + 4HNO

(conc.) → Cu(NO

)

+ 2NO

+ 2H

Zinc reacts with dilute nitric acid to give N

O and with concentrated

acid to give NO

4Zn + 10HNO

(dilute) → 4 Zn (NO

)

+ 5H

O + N

Zn + 4HNO

(conc.) → Zn (NO

)

+ 2H

O + 2NO

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

, and phosphorus to phosphoric acid.

+ 10HNO

→ 2HIO

+ 10NO

+ 4H

C + 4HNO

→ CO

+ 2H

O + 4NO

+ 48HNO

→ 8H

+ 48NO

+ 16H

+ 20HNO

→ 4H

+ 20NO

+ 4H

Brown Ring Test: The familiar brown ring test for nitrates depends

on the ability of Fe

to reduce nitrates to nitric oxide, which reacts

with Fe

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.

+ 3Fe

+ 4H

→ NO + 3Fe

+ 2H

])OH( Fe[

+ NO → [Fe (H

(NO)]

+ H

(brown)

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

(

)

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

molecule where the angles are only 60°. It readily

catches fire in air to give dense white fumes of P

4 2 4 10

P 5O P O

+ →

It consists of discrete tetrahedral P

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

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.

+ 6H

O → 3Ca(OH)

+ 2PH

+ 6HCl → 3CaCl

+ 2PH

In the laboratory, it is prepared by heating white phosphorus with

concentrated NaOH solution in an inert atmosphere of CO

(

)

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

or P

vapours. To purify it from the impurities,

it is absorbed in HI to form phosphonium iodide (PH

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

, Cl

and

vapours.

It is slightly soluble in water. The solution of PH

in water decomposes

in presence of light giving red phosphorus and H

. When absorbed in

P P

Fig.7.3: Red phosphorus

P P

60°

Fig. 7.2

White phosphorus

7.7

Phosphine

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182Chemistry

7.8.1 Phosphorus

Trichloride

Phosphorus forms two types of halides, PX

(X = F, Cl, Br, I) and

(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

COOH, C

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

hybridised.

7.87.8

7.8

Phosphorus HalidesPhosphorus Halides

Phosphorus Halides

UsesUses

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.

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 7.6

In what way can it be proved that PH

is basic in nature?

reacts with acids like HI to form PH

I which shows that it is

basic in nature.

3 4

PH HI PH I

+ →

Due to lone pair on phosphorus atom, PH

is acting as a Lewis base in the above reaction.

SolutionSolution

Solution

Intext QuestionsIntext Questions

Intext Questions

7.7 (a) Bond angle in PH

is higher than that in PH

. Why?

(b) What is formed when PH

reacts with an acid?

7.8 What happens when white phosphorus is heated with concentrated NaOH

solution in an inert atmosphere of CO

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

on phosphorus.

4 2 2 5 2

P 10SO Cl 4PCl 10SO

+ → +

Properties

PCl

is a yellowish white powder and in moist air, it hydrolyses to

POCl

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

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

Cl, CH

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.

240 pm

202 pm

Why does PCl

fume in moisture ?

PCl

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

molecule equivalent? Justify your answer.

PCl

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 7.7

SolutionSolution

Solution

Example 7.8Example 7.8

Example 7.8

SolutionSolution

Solution

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184Chemistry

Intext QuestionsIntext Questions

Intext Questions

7.9 What happens when PCl

is heated?

7.10 Write a balanced equation for the reaction of PCl

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

)

have been given in the Table.

The compositions of the oxoacids are interrelated in terms of loss

or gain of H

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

)

or P–H (e.g., in H

) 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

Oxoacids of

Phosphorus

Hypophosphorous H

+1 One P – OH white P

+ alkali

(Phosphinic) Two P – H

One P = O

Orthophosphorous H

+3 Two P – OH P

+ H

(Phosphonic) One P – H

One P = O

Pyrophosphorous H

+3 Two P – OH PCl

+ H

Two P – H

Two P = O

Hypophosphoric H

+4 Four P – OH red P

+ alkali

Two P = O

One P – P

Orthophosphoric H

+5 Three P – OH P

One P = O

Pyrophosphoric H

+5 Four P – OH heat phosphoric

Two P = O acid

One P – O – P

Metaphosphoric* (HPO

)

+5 Three P – OH phosphorus acid

Three P = O + Br

, heat in a

Three P – O – P sealed tube

Oxidation

state of

phosphorus

Characteristic

bonds and their

number

PreparationName Formula

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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

to metallic silver.

4 AgNO

+ 2H

O + H

→ 4Ag + 4HNO

+ H

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

and

are dibasic and tribasic, respectively as the structure of H

has two P–OH bonds and H

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

7.10

Group 16Group 16

Group 16

ElementsElements

Elements

Fig. 7.4

Structures of some

important oxoacids of

phosphorus

H PO

Orthophosphoric acid

3 4

Orthophosphorous acid

H PO

3 3

H P O

Pyrophosphoric acid

4 2 7

Hypophosphorous acid

H PO

3 2

O O

OH OH

OH O

Cyclotrimetaphosphoric acid, (HPO )

3 3

Polymetaphosphoric acid, (HPO )

3 n

Intext QuestionsIntext Questions

Intext Questions

7.11 What is the basicity of H

7.12 What happens when H

is heated?

How do you account for the reducing behaviour

of H

on the basis of its structure ?

In H

, two H atoms are bonded directly to P

atom which imparts reducing character to the acid.

Example 7.9Example 7.9

Example 7.9

SolutionSolution

Solution

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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

.2H

O, epsom salt MgSO

.7H

O, baryte BaSO

and

sulphides such as galena PbS, zinc blende ZnS, copper pyrites CuFeS

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

. 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

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.

Single bond;

Approximate value;

At the melting point;

Rhombic sulphur;

Hexagonal grey;

Monoclinic form, 673 K.

* Oxygen shows oxidation states of +2 and +1 in oxygen fluorides OF

and O

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

[Ne]3s

[Ar]3d

[Kr]4d

[Xe]4f

Covalent radius/(pm)

66 104 117 137 146

Ionic radius, E

2–

/pm 140 184 198 221 230

Electron gain enthalpy, –141 –200 –195 –190 –174

/∆

H kJ mol

–1

Ionisation enthalpy (∆

) 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

2.06

4.19

6.25 –

Melting point/K 55 393

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

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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

) whereas

sulphur exists as polyatomic molecule (S

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

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

which is not found in H

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 7.10

SolutionSolution

Solution

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188Chemistry

(ii) Reactivity with oxygen: All these elements form oxides of the EO

and EO

types where E = S, Se, Te or Po. Ozone (O

) and sulphur

dioxide (SO

) are gases while selenium dioxide (SeO

) is solid.

Reducing property of dioxide decreases from SO

to TeO

; SO

reducing while TeO

is an oxidising agent. Besides EO

type,

sulphur, selenium and tellurium also form EO

type oxides (SO

SeO

, TeO

). 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

, EX

and EX

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

is exceptionally stable for steric reasons.

Amongst tetrafluorides, SF

is a gas, SeF

a liquid and TeF

a solid.

These fluorides have sp

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

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

E (E = O, S, Se, Te, Po). Some properties of

hydrides are given in Table 7.7. Their acidic character increases

from H

O to H

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

O to H

Po. All the hydrides except

water possess reducing property and this character increases from

S to H

Te.

Property H

O H

S H

Se H

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

∆

H/kJ mol

–1

–286 –20 73 100

∆

diss

H (H–E)/kJ mol

–1

463 347 276

238

Dissociation constant

1.8×10

–16

1.3×10

–7

1.3×10

–4

2.3×10

–3

Aqueous solution, 298 K

Table 7.7: Properties of Hydrides of Group 16 Elements

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189 The p-Block Elements

are S

, S

, Se

and Se

. These dimeric halides undergo

disproportionation as given below:

2Se

→ SeCl

+ 3Se

7.117.11

7.11

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.

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

O(s) → 4Ag(s) + O

(g); 2Pb

(s) → 6PbO(s) + O

(g)

2HgO(s) → 2Hg(l) + O

(g) ; 2PbO

(s) → 2PbO(s) + O

(g)

(iii) Hydrogen peroxide is readily decomposed into water and dioxygen

by catalysts such as finely divided metals and manganese dioxide.

(aq) → 2H

O(1) + O

(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

in 100 cm

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:

O and

O. Molecular oxygen, O

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

S is less acidic than H

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

SolutionSolution

Solution

Intext QuestionsIntext Questions

Intext Questions

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

O a liquid and H

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

) or mixed (Pb

, Fe

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

, Cl

, CO

, N

). For example,

combines with water to give H

, 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.,

, CrO

, V

). The oxides which give a base with water are known

as basic oxides (e.g., Na

O, CaO, BaO). For example, CaO combines

with water to give Ca(OH)

, a base.

(

)

CaO H O Ca

+ →

7.16 Which of the following does not react with oxygen directly?

Zn, Ti, Pt, Fe

7.17 Complete the following reactions:

(i) C

+ O

→

(ii) 4Al + 3 O

→

7.127.12

7.12

SimpleSimple

Simple

OxidesOxides

Oxides

Intext QuestionsIntext Questions

Intext Questions

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:

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

→ 2ZnO + 2SO

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

+ →

CuCl

2 2 2

4HCl O 2Cl 2H O

+ → +

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

reacts with acids as well as alkalies.

( )

[

]

( ) ( )

( )

[ ]

( )

2 62 3 2

2 3 2 3 6

Al(H O)Al O 6HCl 9H O 2 6Cl

aq aq aq

Al O 6NaOH 3H O 2Naaq aq

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

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.

→ 2O

∆H

(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

+ 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

(g) → PbSO

(s) + 4O

(g)

–

(aq) + H

O(l) + O

(g) → 2OH

–

(aq) + I

(s) + O

(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

gas.

7.13

Ozone

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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

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

. It is insoluble in water but dissolves to some

extent in benzene, alcohol and ether. It is readily soluble in CS

Monoclinic sulphur (

ββ

β-sulphur)

Its m.p. is 393 K and specific gravity 1.98. It is soluble in CS

. 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

molecules. These S

molecules are packed to give different crystal structures. The S

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 —

AllotropicAllotropic

Allotropic

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

. It is a resonance

hybrid of two main forms:

7.18 Why does O

act as a powerful oxidising agent?

7.19 How is O

estimated quantitatively?

Intext QuestionsIntext Questions

Intext Questions

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

molecule which has two

unpaired electrons in the antibonding

* orbitals like O

and, hence,

exhibits paramagnetism.

Example 7.12Example 7.12

Example 7.12

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

, the

ring adopts the chair form and

the molecular dimensions are

as shown in Fig. 7.5 (b).

At elevated temperatures

(~1000 K), S

is the dominant

species and is paramagnetic

like O

Preparation

Sulphur dioxide is formed together with a little (6-8%) sulphur trioxide

when sulphur is burnt in air or oxygen:

S(s) + O

(g) → SO

(g)

In the laboratory it is readily generated by treating a sulphite with

dilute sulphuric acid.

(aq) + 2H

(aq) → H

O(l) + SO

(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

→ Na

+ H

O + SO

→ 2NaHSO

In its reaction with water and alkalies, the behaviour of sulphur

dioxide is very similar to that of carbon dioxide.

7.157.15

7.15

SulphurSulphur

Sulphur

DioxideDioxide

Dioxide

(a) (b)

Fig. 7.5: The structures of (a) S

ring in

rhombic sulphur and (b) S

form

2020-21

194Chemistry

Sulphur forms a number of oxoacids such as H

, H

(x = 2 to 5), H

, H

. 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

7.16

Oxoacids ofOxoacids of

Oxoacids of

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

. It is oxidised to

sulphur trioxide by oxygen in the presence of vanadium(V) oxide catalyst.

(g) + Cl

(g) → SO

(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

is angular. It is a resonance hybrid

of the two canonical forms:

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

is used as a

solvent to dissolve a number of organic and inorganic chemicals.

Intext QuestionsIntext Questions

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

molecule. Are

the two S–O bonds in this molecule equal ?

7.22 How is the presence of SO

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

(ii) conversion of SO

to SO

by the reaction with oxygen in the presence

of a catalyst (V

), and

(iii) absorption of SO

in H

to give Oleum (H

A flow diagram for the manufacture of sulphuric acid is shown in

(Fig. 7.7). The SO

produced is purified by removing dust and other

impurities such as arsenic compounds.

The key step in the manufacture of H

is the catalytic oxidation

of SO

with O

to give SO

in the presence of V

(catalyst).

(

)

(

)

(

)

2 5

V O

2 2 3

2SO g O g 2SO g

+ →

196.6 kJmol

−

∆ = −

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

gas from the catalytic converter is

absorbed in concentrated H

to produce oleum. Dilution of oleum

with water gives H

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.

+ H

→ H

(Oleum)

The sulphuric acid obtained by Contact process is 96-98% pure.

7.177.17

7.17

SulphuricSulphuric

Sulphuric

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

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

In aqueous solution, sulphuric acid ionises in two steps.

(aq) + H

O(l) → H

(aq) + HSO

–

(aq);

= very large (

>10)

HSO

–

(aq) + H

O(l) → H

(aq) + SO

(aq) ;

= 1.2 × 10

–2

The larger value of

(

>10) means that H

is largely

dissociated into H

and HSO

–

. Greater the value of dissociation constant

), 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 HX + M

(X = F, Cl, NO

)

(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.

2 4

H SO

→

12C + 11H

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

Cu + 2 H

(conc.) → CuSO

+ SO

+ 2H

S + 2H

(conc.) → 3SO

+ 2H

C + 2H

(conc.) → CO

+ 2 SO

+ 2 H

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

, cryolite Na

AlF

and fluoroapatite 3Ca

(PO

)

.CaF

)

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

.6H

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

. 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

7.18

Group 17Group 17

Group 17

ElementsElements

Elements

Intext QuestionsIntext Questions

Intext Questions

7.23 Mention three areas in which H

plays an important role.

7.24 Write the conditions to maximise the yield of H

by Contact process.

7.25 Why is

K K

a a

2 1

for H

in water ?

What happens when

(i) Concentrated H

is added to calcium fluoride

(ii) SO

is passed through water?

(i) It forms hydrogen fluoride

2 2 4 4

CaF H SO CaSO 2HF

+ → +

(ii) It dissolves SO

to give H

3 2 2 4

SO H O H SO

+ →

SolutionSolution

Solution

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

)

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

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

[Ne]3s

[Ar]3d

[Kr]4d

[Xe]4f

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

4 3.2 3.0 2.7 2.2

∆

Hyd

H(X

–

)/kJ mol

–1

515 381 347 305 –

–

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)

1.66 (203)

3.19(273)

4.94(293)

–

Distance X – X/pm 143

199 228 266 –

Bond dissociation enthalpy 158.8 242.6 192.8 151.1 –

/(kJ mol

–1

)

2.87

1.36 1.09 0.54 –

Table 7.8: Atomic and Physical Properties of Halogens

Radioactive;

Pauling scale;

For the liquid at temperatures (K) given in the parentheses;

solid;

The

half-cell reaction is X

(g) + 2e

–

→

–

(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 7.14

SolutionSolution

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 7.15

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

, has yellow, Cl

, greenish yellow, Br

, red and

, 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

compared to that of Cl

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

molecule

where they are much closer to each other than in case of Cl

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

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.

+ 2X

–

→ 2F

–

+ X

(X = Cl, Br or I)

+ 2X

–

→ 2Cl

–

+ X

(X = Br or I)

+ 2I

–

→ 2Br

–

+ I

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:

( ) ( ) ( )

( )

hyd

diss

1/2

X g X g X g X

– –

∆

→ → →

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

X g H O HX HOXaq aq

where X = Cl or Br

4I 4H O g 2I 2H Oaq aq

+ −

− +

+ → + +

+ → +

+ + → +

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

/kJ mol

–1

574 432 363 295

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

and

. However, only OF

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

oxidises plutonium to PuF

and the reaction is used

in removing plutonium as PuF

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

O, ClO

, Cl

and Cl

are highly reactive

oxidising agents and tend to explode. ClO

is used as a bleaching

agent for paper pulp and textiles and in water treatment.

The bromine oxides, Br

O, BrO

, BrO

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

, I

are insoluble solids and

decompose on heating. I

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

, PbCl

, SbCl

and UF

are more covalent than

SnCl

, PbCl

, SbCl

and UF

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

′

, XX

′

and XX

′

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 7.16

SolutionSolution

Solution

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

and Cl

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

. 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

+ 4HCl → MnCl

+ Cl

+ 2H

However, a mixture of common salt and concentrated H

used in place of HCl.

4NaCl + MnO

+ 4H

→ MnCl

+ 4NaHSO

+ 2H

O + Cl

(ii) By the action of HCl on potassium permanganate.

2KMnO

+ 16HCl → 2KCl + 2MnCl

+ 8H

O + 5Cl

Manufacture of chlorine

(i) Deacon’s process: By oxidation of hydrogen chloride gas by

atmospheric oxygen in the presence of CuCl

(catalyst) at 723 K.

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

7.19

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

→ 2AlCl

; P

+ 6Cl

→ 4PCl

2Na + Cl

→ 2NaCl; S

+ 4Cl

→ 4S

2Fe + 3Cl

→ 2FeCl

;

It has great affinity for hydrogen. It reacts with compounds

containing hydrogen to form HCl.

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

→ NaCl + NaOCl + H

(cold and dilute)

6 NaOH + 3Cl

→ 5NaCl + NaClO

+ 3H

(hot and conc.)

With dry slaked lime it gives bleaching powder.

2Ca(OH)

+ 2Cl

→ Ca(OCl)

+ CaCl

+ 2H

The composition of bleaching powder is Ca(OCl)

.CaCl

.Ca(OH)

.2H

Chlorine reacts with hydrocarbons and gives substitution products

with saturated hydrocarbons and addition products with unsaturated

hydrocarbons. For example,

+ Cl

→

Cl + HCl

Methane Methyl chloride

+ Cl

Room temp.

→

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

+ H

+ Cl

→ Fe

(SO

)

+ 2HCl

+ Cl

+ H

O → Na

+ 2HCl

+ 2H

O + Cl

→ H

+ 2HCl

+ 6H

O + 5Cl

→ 2HIO

+ 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.

+ H

O → 2HCl + O

Coloured substance + O → Colourless substance

2020-21

204Chemistry

Write the balanced chemical equation for the reaction of Cl

with hot

and concentrated NaOH. Is this reaction a disproportionation

reaction? Justify.

3Cl

+ 6NaOH → 5NaCl + NaClO

+ 3H

Yes, chlorine from zero oxidation state is changed to –1 and +5

oxidation states.

Example 7.17Example 7.17

Example 7.17

SolutionSolution

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

420 K

→

NaHSO

+ HCl

NaHSO

+ NaCl

823K

→

+ 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:

(

)

( )

(

)

(

)

2 3 a

HCl g H O H O Cl 10

aq aq

+ −

+ → + =K

Its aqueous solution is called hydrochloric acid. High value of

dissociation constant (K

) indicates that it is a strong acid in water.

It reacts with NH

and gives white fumes of NH

Cl.

+ HCl → NH

When three parts of concentrated HCl and one part of concentrated

HNO

are mixed, aqua regia is formed which is used for dissolving

noble metals, e.g., gold, platinum.

3 4 2

3 6 2

Au 4H NO 4Cl AuCl NO 2H O

3Pt 16H 4NO 18Cl 3PtCl 4NO 8H O

+ − − −

+ + + → + +

7.207.20

7.20

HydrogenHydrogen

Hydrogen

ChlorideChloride

Chloride

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

, CHCl

, DDT, refrigerants, etc. (iv) in sterilising drinking water

and (v) preparation of poisonous gases such as phosgene (COCl

), tear gas

(CCl

), mustard gas (ClCH

SCH

Cl).

Intext QuestionsIntext Questions

Intext Questions

7.29 Give the reason for bleaching action of Cl

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.

+ 2HCl → 2NaCl + H

O + CO

NaHCO

+ HCl → NaCl + H

O + CO

+ 2HCl → 2NaCl + H

O + SO

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

HOBrO

HOIO

(Halic acid) – (chloric acid) (bromic acid) (iodic acid)

Halic (VII) acid – HOCIO

HOBrO

HOIO

(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

Fe 2HCl FeCl H

+ → +

Liberation of hydrogen prevents the formation of ferric chloride.

Example 7.18Example 7.18

Example 7.18

SolutionSolution

Solution

UsesUses

Uses: It is used (i) in the manufacture of chlorine, NH

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 of

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

′

and XX

′

where X is halogen of larger size and X

′

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

(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

7.22

InterhalogenInterhalogen

Interhalogen

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′

ClF colourless gas –

BrF pale brown gas –

detected spectroscopically –

BrCl

gas

ICl ruby red solid (α-form) –

brown red solid (β-form) –

IBr black solid –

XX′

ClF

colourless gas Bent T-shaped

BrF

yellow green liquid Bent T-shaped

yellow powder Bent T-shaped (?)

ICl

orange solid Bent T-shaped (?)

XX′

colourless gas but Square

solid below 77 K pyramidal

BrF

colourless liquid Square

pyramidal

ClF

colourless liquid Square

pyramidal

XX′

colourless gas Pentagonal

bipyramidal

Table 7.11: Some Properties of Interhalogen Compounds

Very unstable;

The pure solid is known at room temperature;

Dimerises as Cl–bridged

dimer (I

)

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′

), halate (when XX′

) and perhalate (when XX′

) anion derived

from the larger halogen.

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

compounds have the bent ‘T’ shape, XX

compounds square pyramidal

and IF

has pentagonal bipyramidal structures (Table 7.11).

Discuss the molecular shape of BrF

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 7.19

SolutionSolution

Solution

UsesUses

Uses: These compounds can be used as non aqueous solvents. Interhalogen

compounds are very useful fluorinating agents. ClF

and BrF

are used for the

production of UF

in the enrichment of

235

U(s) + 3ClF

(l) → UF

(g) + 3ClF(g)

Intext QuestionIntext Question

Intext Question

7.31 Why is ICl more reactive than I

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

atoms and

ions

249 48 294

98 20 118

Cf Ca Og + 3

+ n

7.237.23

7.23

Group 18Group 18

Group 18

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 7.20

SolutionSolution

Solution

Oganesson has its symbol Og, atomic number 118, atomic mass

294 and electronic configuration [Rn] 5f

. 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

[He]2s

[Ne] 3s

[Ar]3d

[Kr]4d

[Xe]4f

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

except

helium which has 1s

(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

) have completely filled ns

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

PtF

–

. 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

–

by mixing PtF

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

) 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 7.21

SolutionSolution

Solution

2020-21

210Chemistry

Fig. 7.9

The structures of

(a) XeF

(b) XeF

(d) XeOF

and (e) XeO

(a) Linear

(b) Square planar

(e) Pyramidal

but only identified (e.g., RnF

) by radiotracer technique. No true

compounds of Ar, Ne or He are yet known.

(a) Xenon-fluorine compounds

Xenon forms three binary fluorides, XeF

, XeF

and XeF

by the

direct reaction of elements under appropriate experimental conditions.

Xe (g) + F

(g)

673 K, 1 bar

→

XeF

(s)

(xenon in excess)

Xe (g) + 2F

(g)

873 K, 7 bar

→

XeF

(s)

(1:5 ratio)

Xe (g) + 3F

(g)

573 K , 60 70bar−

→

XeF

(s)

(1:20 ratio)

XeF

can also be prepared by the interaction of XeF

and O

at 143K.

4 2 2 6 2

XeF O F XeF O

+ → +

XeF

, XeF

and XeF

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

is hydrolysed to

give Xe, HF and O

2XeF

(s) + 2H

O(l) → 2Xe (g) + 4 HF(aq) + O

(g)

The structures of the three xenon fluorides can be deduced from

VSEPR and these are shown in Fig. 7.9. XeF

and XeF

have linear and

square planar structures respectively. XeF

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

+ PF

→ [XeF]

[PF

]

–

; XeF

+ SbF

→ [XeF

]

[SbF

]

–

XeF

+ MF → M

[XeF

]

–

(M = Na, K, Rb or Cs)

(b) Xenon-oxygen compounds

Hydrolysis of XeF

and XeF

with water gives Xe0

6XeF

+ 12 H

O → 4Xe + 2Xe0

+ 24 HF + 3 O

XeF

+ 3 H

O → XeO

+ 6 HF

Partial hydrolysis of XeF

gives oxyfluorides, XeOF

and XeO

XeF

+ H

O → XeOF

+ 2 HF

XeF

+ 2 H

O → XeO

+ 4HF

2020-21

211 The p-Block Elements

XeO

is a colourless explosive solid and has a pyramidal molecular

structure (Fig. 7.9). XeOF

is a colourless volatile liquid and has a

square pyramidal molecular structure (Fig.7.9).

Does the hydrolysis of XeF

lead to a redox reaction?

No, the products of hydrolysis are XeOF

and XeO

where the oxidation

states of all the elements remain the same as it was in the reacting state.

Example 7.22Example 7.22

Example 7.22

SolutionSolution

Solution

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 Questions

7.32 Why is helium used in diving apparatus?

7.33 Balance the following equation: XeF

+ H

O → XeO

+ HF

7.34 Why has it been difficult to study the chemistry of radon?

2020-21

212Chemistry

SummarySummary

Summary

Groups 13 to 18 of the periodic table consist of p-block elements with their valence

shell electronic configuration ns

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

. Nitrogen differs from other elements of this

group due to small size, formation of p

ππ

–p

ππ

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

O, NO, N

, NO

, N

and N

. These

oxides have resonating structures and have multiple bonds. Ammonia can be

prepared on large scale by Haber’s process. HNO

is an important industrial

chemical. It is a strong monobasic acid and is a powerful oxidising agent. Metals

and non-metals react with HNO

under different conditions to give NO or NO

Phosphorus exists as P

in elemental form. It exists in several allotropic forms.

It forms hydride, PH

which is a highly poisonous gas. It forms two types of halides as

and PX

. PCl

is prepared by the reaction of white phosphorus with dry chlorine

while PCl

is prepared by the reaction of phosphorus with SO

. 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

. 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

in presence of MnO

. It forms a number of oxides with metals. Allotropic form

of oxygen is O

which is a highly oxidising agent. Sulphur forms a number of allotropes.

Of these,

– and

– forms of sulphur are the most important. Sulphur combines with

oxygen to give oxides such as SO

and SO

. SO

is prepared by the direct union of

sulphur with oxygen. SO

is used in the manufacture of H

. Sulphur forms a

number of oxoacids. Amongst them, the most important is H

. 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

HCl is prepared by heating NaCl with concentrated H

. Halogens combine with

one another to form interhalogen compounds of the type XX

(n = 1, 3, 5, 7)

where X

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

valence

shell electronic configuration except He which has 1s

. 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

form hydrogen bond but PH

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

7.8 Give the resonating structures of NO

and N

7.9 The HNH angle value is higher than HPH, HAsH and HSbH angles. Why?

[Hint: Can be explained on the basis of sp

hybridisation in NH

and only

s–p bonding between hydrogen and other elements of the group].

7.10 Why does R

P = O exist but R

N = O does not (R = alkyl group)?

7.11 Explain why NH

is basic while BiH

is only feebly basic.

7.12 Nitrogen exists as diatomic molecule and phosphorus as P

. 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

7.16 Can PCl

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

by contact process?

7.22 How is SO

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

7.27 Why are halogens coloured?

7.28 Write the reactions of F

and Cl

with water.

7.29 How can you prepare Cl

from HCl and HCl from Cl

? Write reactions only.

7.30 What inspired N. Bartlett for carrying out reaction between Xe and PtF

7.31 What are the oxidation states of phosphorus in the following:

(i) H

(ii) PCl

(iii) Ca

(iv) Na

(v) POF

Exercises

2020-21

214Chemistry

7.32 Write balanced equations for the following:

(i) NaCl is heated with sulphuric acid in the presence of MnO

(ii) Chlorine gas is passed into a solution of NaI in water.

7.33 How are xenon fluorides XeF

, XeF

and XeF

obtained?

7.34 With what neutral molecule is ClO

–

isoelectronic? Is that molecule a Lewis

base?

7.35 How are XeO

and XeOF

prepared?

7.36 Arrange the following in the order of property indicated for each set:

(i) F

, Cl

, Br

, I

- increasing bond dissociation enthalpy.

(ii) HF, HCl, HBr, HI - increasing acid strength.

(iii) NH

, PH

, AsH

, SbH

, BiH

– increasing base strength.

7.37 Which one of the following does not exist?

(i) XeOF

(ii) NeF

(iii) XeF

(iv) XeF

7.38 Give the formula and describe the structure of a noble gas species which

is isostructural with:

(i) ICl

–

(ii) IBr

–

(iii) BrO

–

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

is the least stable among the hydrides of Group 15.

7.3 Because of strong pπ–pπ overlap resulting into the triple bond, N≡N.

7.6 From the structure of N

it is evident that covalence of nitrogen is four.

7.7(a) Both are sp

hybridised. In PH

all the four orbitals are bonded whereas

in PH

there is a lone pair of electrons on P, which is responsible for lone

pair-bond pair repulsion in PH

reducing the bond angle to less than

109° 28′.

7.10 PCl

+ H

O → POCl

+ 2HCl

7.11 Three P–OH groups are present in the molecule of H

. 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

is a very strong acid in water largely because of its first ionisation

to H

and HSO

–

. The ionisation of HSO

–

to H

and SO

2–

is very very

small. That is why K

<< K

7.31 In general, interhalogen compounds are more reactive than halogens due

to weaker X–X

bonding than X–X bond. Thus, ICl is more reactive than I

7.34 Radon is radioactive with very short half-life which makes the study of

chemistry of radon difficult.

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