Chemistry

c. Group 18 elements (Noble gases) : iv. Nature of hydrides : Hydride of oxygen
(H2O) is liquid at room temperature while
Noble gases are monoatomic. hydrides of other members of the group are
gases.
They are sparingly soluble in water.
Use your brain power
Noble gases have very low melting and • Which of the following possess
boiling points. Helium has the lowest boiling
point (4.2 K) of any known substances. hydrogen bonding ? H2S, H2O,
H2Se, H2Te
Problem 7.6 : Noble gases have very low • Show hydrogen bonding in the above
melting and boiling points. Why ? molecule with the help of a diagram.

Solution : Noble gases are monoatomic, v. Common covalency of oxygen is 2. In rare
the only type of inter atomic interactions cases it is four. But for the other members of
which exist between them are van der the group 16 the covalency can exceed four.
Waals forces. Therefore, they can be
liquified at very low temperatures and The anomalous behaviour of oxygen is due
have very low melting or boiling points. to the following reasons.

Can you tell ? i. small atomic size
The first member of a group
usually differs in properties from ii. high electronegativity.
the rest of the members of the group.
Why ? iii. absence of inner d-orbitals.

7.5 Anamalous Behaviour 7.5.2 Anomalous behaviour of fluorine :
Fluorine, the first member of group 17, differs
7.5.1 Anomalous behaviour of oxygen : in properties from the other members of the
Oxygen shows the following anomalous group. The anomalous behaviour of fluorine
properties compared to other members of is due to the following reasons.
group 16 :
i. small atomic size
i. Atomicity : Oxygen is a diatomic molecule
(O2) while others are polyatomic molecules. ii. high electronegativity
For example P4, S8.
ii. Magnetic property : Oxygen is iii. absence of d-orbitals in valence shell
paramagnetic while others are diamagnetic.
iv. low F-F bond dissociation enthalpy
iii. Oxidation state : Oxygen shows -2, -1,
and +2 oxidation states while other elements Some anomalous properties of fluorine :
show, -2, +2, +4, +6 oxidation states. Oxygen
can not exhibit higher oxidation state due to i. Ionisation enthalpy, electronegativity,
absence of vacant d orbitals. electrode potential are all higher for fluorine
than expected trends shown by other halogens.
Use your brain power
• Oxygen forms only OF2 with ii. Ionic and covalent radii, m.p., b.p. and
electron gain enthalpy are quite lower than
fluorine while sulfur forms SF6. expected.
Explain. Why ?
iii. Most of the reactions of fluorine are
exothermic (due to the short and strong bond
formed by it with other elements).

142

iv. It forms only one oxoacid (HOF) while All halogens exhibit -1 oxidation state.
other halogens form a number of oxoacids. However Cl, Br and I exhibit +1, +3, +5
and +7 oxidation states as well. This is
v. Hydrogen fluoride is a liquid (b.p. 293K) because they are less electronegative than F
due to strong hydrogen bonding while other and possess empty d-orbitals in the valence
hydrogen halides are gases. shell and therefore, can expand the octet. The
oxidation states +4 and +6 occur in the oxides
7.6 Chemical Properties of elements of and oxoacids of Cl and Br.
groups 16, 17 and 18
The fluorine atom has no d - orbitals
7.6.1 Oxidation state : i. The group 16 in its valence shell and therefore cannot
elements have the valence shell electronic expand its octet. Thus fluorine being most
configuration ns2np4. They attain a noble gas electronegative exhibits mostly -1 oxidation
configuration either by gaining two electrons, state.
forming E2 ions or by sharing two electrons,
forming two covalent bonds. These elements, iii. Group 18 elements (noble gases)
thus, show -2 and +2 oxidation states in their have stable valence shell electronic
compounds. configuration ns2np6 with completely filled
orbitals. Thus they have no tendency to
Oxygen being highly electronegative, gain or lose electrons, that is, they are zero
shows common oxidation state of -2 except valent and mostly exist as monoatomic gases.
two cases. In the case of OF2, its oxidation However, xenon has large atomic size and
state is +2 and in peroxides, it shows oxidation lower ionisation enthalpy compared to He,
state -1 (H2O2, Na2O2). Other elements of the Ne, Ar and Kr. Hence xenon exhibits higher
group exhibit +2, +4, +6 oxidation states with oxidation states. Its outermost shell has
+4 and +6 being more common. The stability d-orbitals. The paired electrons of the valence
of higher (+6) oxidation state decreases down shell can be unpaired and promoted to empty
the group while the stability of the lower d-orbitals. The unpaired electrons are shared
oxidation state (+4) increases down the group with fluorine or oxygen atoms and covalent
due to inert pair effect. Bonding in +4 and +6 compounds showing higher oxidation state
oxidation states are primarily covalent. such as XeF2(+2), XeF4 (+4), XeF6 (+6), XeO3
(+6) and XeOF4 (+6) are formed.
Try this...
Complete the following tables 7.6.2 Chemical Reactivity towards
hydrogen:
Element O OS F
H2O OF2 H2S HF i. Group 16 elements : The elements of group
compound -2 16 react with hydrogen to form hydrides of
Se Te Cl the type H2E. (Where E = O, S, Se, Te, Po).
Oxidation Se TeF6 HOCl
state SeO2 SeO3 For example, H2O, H2S, H2Se, H2Te and H2Po.
Element +6 Some properties of hydrides of group 16 are
compound given in Table 7.5.

Oxidation
state

ii. The group 17 elements are represented H2O is a colourless, odourless liquid, while
by their valence shell electronic configuration H2S, H2Se, H2Te and H2Po are colourless
as ns2np5. They attain noble gas configuration bad smelling, poisonous gases at ambient
either by gaining one electron forming E conditions.
ions or by sharing one electron forming one
covalent bond. All hydrides have angular structures which
involve sp3 hybridisation of central atom (E).

143

The hydrides of group 16 elements are weakly Table 7.6 Properties of hydrides of group
acidic. The acidic character of the hydrides 17 elements.
increases, while thermal stability decreases
from H2O to H2Te. This is due to decrease Property HF HCl HBr HI
in the bond dissociation enthalpy of the H-E
bond down the group (Table 7.5). m.p (K) 190 159 185 222

b.p (K) 293 189 206 238

All hydrides except H2O possess reducing Bond length 91.7 127.4 141.4 160.9
property which increases in the order H2S < (H-X) pm
H2Se < H2Te.
∆dissH0 574 432 363 295
ii. Group 17 elements : The elements of kJ/mol
group 17 react with hydrogen to give hydrogen
halides. pKa 3.2 -7.0 -9.5 -10.0

iii. Group 18 elements (Noble gases) :

H2 + X2 2HX Noble gases are chemically inert towards
hydrogen due to their stable electronic
(Where X = F, Cl, Br, I) configuration.

Some of the properties of hydrogen halides 7.6.3 Reactivity towards oxygen :
are given in Table 7.6.
i. Group 16 elements : All the elements of
Acidic strength of halogen acids increases in group 16 form oxides of the type EO2 and
the order : EO3 where E = S, Se, Te, Po.

HF < HCl < HBr < HI EO2 type oxides, Ozone (O3) and sulfur
dioxide (SO2) are gases, while selenium
It is due to decreasing bond dissociation dioxide (SeO2) is solid. They are acidic in
enthalpy of H-X bond in the order nature and react with water to form acids.
HF > HCl > HBr > HI.

Thermal stability of hydrogen halides SO2 + H2O H2SO3 (Sulfurous acid)
decreases in the order HF>HCl>HBr>HI. It is
due to decrease in bond dissociation enthalpy SeO2 + H2O H2SeO3 (Selenious acid)
of H-X bond down the group.
Reducing property of dioxides decreases from
Table 7.5 Properties of hydrides of group
16 elements SO2 to TeO2. SO2 is reducing while TeO2
serves as an oxidising agent.

Property H2O H2S H2Se H2Te EO3 type oxides, SO3, SeO3, TeO3 are also
m.p (K) 273 188 208 222 acidic in nature. They dissolve in water to
form acids.

b.p (K) 373 213 232 269 SeO3 + H2O H2SeO4 (Selenic acid)

H-E TeO3 + 3H2O H6TeO6 (Telluric acid)
Bond length
96 134 146 169 ii. Group 17 elements : Elements of group 17
(pm)
(Halogens) form many oxides with oxygen,
∆dissH (H-E)
kJ/mol 463 347 276 238 but most of them are unstable.

∆fH kJ/mol -286 -20 73 100 Fluorine forms two oxides OF2 and O2F2.
However, only the OF2 is thermally stable at
HEH angle (°) 104 92 91 90 298 K. Both are strong fluorinating agents.

pKa 14.0 7.0 3.8 2.6 O2F2 oxidises plutonium to PuF6 and the
reaction is used in removing plutonium as

PuF6 from spent nuclear fuel.

144

Chlorine oxides, Cl2O, ClO2, Cl2O6 and Cl2O7 Internet my friend
are highly reactive oxidising agents and tend
to explode. Find and draw the structures of
SeF4 and SCl2.
ClO2 is used as bleaching agent for paper
pulp and textiles and in water treatment. ii. Group 17 elements :

Bromine oxides, Br2O, BrO2, BrO3 are the Halogens (Group 17 elements) combine
least stable halogen oxides (middle row amongst themselves to form a number of
anomally). They are very powerful oxidising compounds known as interhalogen compounds.
agents.
These are of following types : XX', XX'3,
Iodine oxides, I2O4, I2O5 and I2O7 are insoluble XX'5, XX'7,
solids and decompose on heating. I2O5 is a
very good oxidising agent and used for the Where X is the halogen atom with larger size
estimation of carbon monoxide. and X', is the halogen atom with smaller size.
More details of interhalogen compounds are
The higher oxides of halogens are more stable included in section 7.12.
than the lower ones.
iii. Group 18 elements : Group 18 elements
iii. Group 18 elements : Noble gas elements
are chemically inert and do not directly react (Noble gases) are chemically inert. Krypton
with oxygen.
and xenon, however react directly with
7.6.4 Reactivity towards halogens :
fluorine to give their fluorides. For example,
i. Group 16 elements : Elements of group 16
react with halogens to give a large number Xe(g) + F2(g) 673K XeF2(s)
of halides of the types EX6 , EX4 and EX2. 1 atm
(Where E = S, Se, Te)
Xenon fluorides XeF2, XeF4 and XeF6 are
Hexahalides, SF6, SeF6 and TeF6 are formed by crystalline and colourless which sublime
direct combination. They are colourless gases.
They have sp3d2 hybridisation and possess readily at 298 K. They are powerful
octahedral structure. SF6 is exceptionally
stable halide for steric reasons. fluorinating agents.

Stability of halides decreases in the order 7.6.5 Reactivity towards metals :
fluorides > chlorides > bromides > iodides
i. Group 16 elements : Elements of group
16 react with metals to form corresponding
compounds.

e.g. 4Al + 3O2 2Al2O3
Cu + S CuS

Tetrahalides, SF4, SeF4, TeF4, TeCl4 have sp3 Mg + Se MgSe
hybridisation and thus trigonal bipyramidal
geometry with one equatorial position magnesium selenide
occupied by a lone pair.
Do you know ?
Dihalides, SCl2, SeCl2, TeCl2 have sp3
hybridisation and thus possess tetrahedral Tellurium has the unusal property
structure with two equatorial positions of combining with gold metal to
occupied by lone pairs. form telluride.

Monohalides are dimeric in nature. For 2Au + 3Te Au2Te3
example, S2F2, S2Cl2, Se2Cl2 and SeBr2. These (gold telluride)
dimeric halides undergo disproportionation.

2 Se2Cl2 SeCl4 + 3Se

145

Internet my friend Do you know ?
www.chemistry.explained.com The photocopying process.

ii. Group 17 elements : A selenium-coated rotating
drum is given a uniform positive charge
Elements of group 17 (Halogens) react with (step 1) and is then exposed to an image
metals instantly to give metal halides. (step 2). Negatively charged toner particles
are attracted to the charged area of the
2Na (s) + Cl2 (l) 2NaCl (s) drum (step 3) and the image is transferred
Mg (s) + Br2 (l) MgBr2 (s) from the drum to a sheet of paper (step
magnesium bromide 4). Heating then fixes the image and the
drum is flooded with light and cleaned to
Ionic character of halides decreases in the ready the machine for another cycle (step
order MF > MCl > MBr > MI, where M is a 5). Figure of photocopying process using
monovalent metal. Se is as shown below :

The metal halides having metals in their higher • Tellurium exists in two allotropic forms (i)
oxidation states are more covalent than the crystalline and (ii) amorphous.
ones having metals in lower oxidation state.
For example, SnCl4, PbCl4, SbCl5 and UF6 are • Polonium reveals two allotropic forms α
more covalent than SnCl2, PbCl2, SbCl3 and and β (both metallic).
UF4 respectively.
7.7.1 Allotropes of sulfur :
iii. Group 18 elements : Noble gases do not Sulfur exhibits numerous allotropic
directly react with metals.
forms. However rhombic sulfur (α- sulfur)
7.7 Allotropy : and monoclinic sulfur (β - sulfur) are the most
important allotropes of sulfur (Table 7.7).
Can you tell ?
• What is allotropy ? Problem 7.7 : Which form of sulfur shows
paramagnetic behaviour ?
• What is the difference between Solution : In vapour state, sulfur partly
allotropy and polymorphism ? exists as S2 molecule, which has two
unpaired electrons in the antibonding
Elements of the group 16 exhibit π* orbitals like O2. Hence it exhibits
allotropy. Oxygen has two allotropes O2 and paramagnetism.
O3 (ozone). Sulfur exists in a number of
allotropic forms. Rhombic and monoclinic
sulfur are the important allotropes of sulfur.
Both are non metallic.

• Selenium exists in two allotropic forms red
(non metallic) and grey (metallic).

Do you know ?

Grey selenium allotrope of is a
photoconductor used in photocells.

146

Table 7.7 Allotropes of sulfur

Rhombic Sulfur (α - Sulfur) Monoclinic Sulfur (β - Sulfur)

Color Pale Yellow Bright yellow solid
Shape
M. P. Orthorhombic crystals Needle shaped monoclinic crystals
Density
Solubility 385.8 K 393 K

Stability 2.069/ cm3 1.989 / cm3

Structure Insoluble in water and soluble in Soluble in CS2
CS2
Method of
preparation Stable below 369 K and transforms Stable above 369 K and transforms

to β - Sulphur above this into α - Sulphur below this

temperature. temperature.

S8 molecules having puckered ring S8 molecules with puckered ring
structure structure

It is prepared by evaporation of Rhomic sulphur melted in a dish
roll sulphur in CS2. and cooled till crust is formed. Two
holes are made in the crust and
remaining liquid is poured out to
give needle shaped crystals of
β - Sulphur

Remember... 7.8 Oxoacids

Several modifications of sulfur 7.8.1 Oxoacids of sulfur : Sulfur forms
containing 6-20 sulfur atoms per ring, have a number of oxoacids. Some of them are
been synsthesised. In the S8 molecule the unstable and cannot be isolated. They are
ring is puckered and has a crown shape. known to exist in aqueous solutions or in the
In cyclo - S6, the ring adopts the chair form of their salts.
form. At elevated temperature (~ 1000 K),
S2 is the dominant species which like O2 Some important oxoacids of sulfur and
is paramagnetic. their structures are given below.

i. Sulfurous acid, H2SO3

S 2.04 S S HO S O
pm S 107° HO
S

S ii. Sulfuric Acid, H2SO4
SS O

Structure of S8 ring in rhombic sulfur HO S O
HO
S
S 205.7°c S iii. Di or pyrosulfuric acid, H2S2O7

S 02.2° S OO

S O S OO S
OH O
Chair form of S6
HO

147

iv. Peroxy monosulfuric acid, H2SO5 Strucutres of oxoacids of chlorine :
O
i. Hypochlorous acid, HOCl
HOO S O
HO O
H Cl
v. Peroxy disulfuric acid, H2S2O8 ii. Chlorous acid, HOClO or HClO2

OO
H Cl

SS iii. Chloric acid, HClO3
OO
OH HO O O O
H Cl O
vi. Thiosulfuric acid, H2S2O3 iv. Perchloric acid, HClO4

S

HO S O O
HO O
H Cl O

O

7.8.2 Oxoacids of halogens : :Halogens form 7.9 Oxygen and Compounds of oxygen
several oxoacids (See Table 7.8). Only four
oxoacids have been isolated in pure form: 7.9.1 Dioxygen
hypofluorous acid (HOF), perchloric acid
(HClO4), iodic acid (HIO3), metaperiodic acid a. Preparation
(H2IO6). The others are stable only in aqueous
solutions or in the form of their salts. i. Laboratory methods :

The acid strength of the halogen oxoacids • By heating oxygen containing salts such
increases with the increasing oxidation state of as chlorates, nitrates and permanganates.
halogen. For example, acid strength increases
from HClO, a weak acid (Ka = 3.5 × 10-8), to 2KClO3(s) Heat 2KCl(s) + 3O2(g)
HClO4, a very strong acid (Ka>>1). MnO2
• By thermal decomposition of oxides of
metals.

2Ag2O(s) ∆ 4Ag(s) + O2(g)
2HgO(s) ∆ 2Hg(l) + O2(g)

Table 7.8 Oxoacids of halogens

oxidation state Generic Oxoacids of Oxoacids of Oxoacids of Oxoacids of
of X name bromine iodine
+1 fluorine chlorine HOBr HOI

+3 Hypohalous HOF HOCl - -
acid (HXO)
+5 HOBrO2 HOIO2
Halous acid - HOClO
+7 (HXO2) HOBrO3 HOIO3

Halic acid - HOClO2
- HOClO3
(HXO4)
Perhalic acid

(HXO4)

148

2PbO2(s) ∆ 2PbO(s) + O2(g) iii. Reaction with some compounds :

• By decomposition of hydrogen peroxide in 2ZnS + 3O2 ∆ 2ZnO + 2SO2

presence of catalyst such as finely divided CH4 + 2O2 CO2 + 2H2O

metals and manganese dioxide. 2SO2 + O2 V2O5 2SO3

2H2O2(aq) ∆ 2H2O(l) + O2(g) 4HCl + O2 CuCl2 2Cl2 + 2H2O

MnO2 ∆
ii. Electrolysis : Dioxygen can be prepared
d. Uses
on large scale by electrolysis of water, when
• Dioxygen is important for respiration to
hydrogen is liberated at cathode and oxygen sustain animal and aquatic life.

at anode. • It is used in the manufacture of steel.

2H2O Electrolysis 2H2 + O2 • It is used in oxyacetylene flame for welding
and cutting of metals.
iii. Industrial method : Dioxygen is obtained

from air, by first removing carbon dioxide • Oxygen cylinders are widely used
in hospitals, high altitude flying and
and water vapour. The remaining gases are mountaineering.

liquified subsequently. This is followed by

fractional distillation which gives dinitrogen

and dioxygen. • It is used in combustion of fuels; for
example, hydrazine in liquid oxygen
b. Physical properties : provides tremendous thrust (energy) in
• Dioxygen is colourless and odourless gas. rockets.

• Dioxygen is sparingly soluble in water, 30.8 Try this...
cm3 of O2 dissolves in 1000 cm3 of water
at 293 K. A small amount of dissolved Why water in the fish pot needs
dioxygen is sufficient to sustain marine and to be changed time to time ?
aquatic life.

• It liquifies at 90 K and freezes at 55 K. Problem : 7.8

• Oxygen has three stable isotopes 16O, 17O Dioxygen is paramagnetic inspite of having
and 18O. even number of electrons. Explain.

• Molecular oxygen, O2 exhibits Solution : Dioxygen is a covalently bonded
paramagnetism. molecule.

c. Chemcial Properties : O O or O O
i. Reaction with metals : Dioxygen directly Paramagnetic behaviour of O2 can be
reacts with almost all metals except Au, Pt explained with the help of molecular
to form their oxides. orbital theory .

2Ca + O2 2CaO Electronic configuration of O2
KK σ(2s)2 σ*(2s)2 σ(2pz)2 π(2px)2 π(2py)2
4Al + 3O2 2Al2O3 π*(2px)1 π*(2py)1. Presence of two unpaired
electrons explains paramagnetic nature of
ii. Reaction with nonmetals : Dioxygen dioxygen.

reacts with nonmetals (except noble gases)

to form their oxides.

C + O2 CO2

P4 + 5O2 P4O10

149

7.9.2 Simple Oxides : A binary compound The layer of ozone protects the earth’s surface
of oxygen with another element is called an from harmful ultraviolet (U.V) radiations.
oxide. Hence, it is called as 'ozone umbrella'.

Oxides can be classified into a. Preparation of Ozone : Ozone is prepared

a. Acidic oxides b. Basic oxides in the laboratory by passing silent electric
c. Amphoteric oxides d. Neutral oxides discharge through pure and dry oxygen in an
apparatus called ozoniser. As the conversion

a. Acidic oxides : An oxide which dissolves of oxygen to ozone is only 10%, the product
in water to give an acid or reacts with a is known as ozonised oxygen. It is an
base to give a salt is called acidic oxide. For endothermic process.

example, SO2, SO3, CO2, N2O5, Cl2O7 etc. 3O2(g) 2O3 ∆H = +142 kJ/mol

SO2 + H2O H2SO3 b. Physical properties of ozone :

SO3 + 2NaOH Na2SO4 + H2O i. Pure Ozone is a pale - blue gas, dark blue
liquid and violet - black solid.
Generally, oxides of nonmetals are acidic

oxides. ii. Ozone has a characteristic smell. When

b. Basic oxides : An oxide which dissolves in inhaled in concerntration above 100 ppm, it
water to give a base or reacts with an acid to causes nausea and headache.

give salt is called basic oxide. For example, iii. It is diamagnetic in nature.

Na2O, CaO, BaO etc. Ca(OH)2 Problem 7.9 : High concerntration of ozone
CaO + H2O can be dangerously explosive. Explain.

BaO + 2HCl BaCl2 + H2O Solution : i. Thermal stability : Ozone is
thermodynamically unstable than oxygen
c. Amphoteric oxides : The oxide which and decomposes into O2. The decomposition
reacts with a base as well as with an acid to is exothermic and results in the liberation
give salt is called an amphoteric oxide. For of heat (∆H is -ve) and an increase in
example, Al2O3 entropy (∆S is positive). This results in
large negative Gibbs energy change (∆G).
Al2O3 + 6NaOH(aq) + 3H2O(l) 2Na3[Al(OH)6] Therefore high concerntration of ozone can

(Acidic)

Al2O3(s)+ 6HCl(aq)+ 9H2O(l) 2[Al(H2O)6]3⊕ be dangerously explosive.

(Basic) Eq. O3 O2 + O

+ 6Cl (aq) c. Chemical Properties :

d. Neutral oxides : The oxides which are i. Oxidising property :

neither acidic nor basic, are called as neutral

oxides. For example, CO, NO, N2O etc. Ozone is a powerful oxidising agents as it
easily decomposes to liberate nascent oxygen.
7.9.3 Ozone : Ozone (O3) is an allotrope (O3 O2 + O).
of oxygen. Oxygen in the upper atmosphere
Ozone oxidises lead sulfide to lead sulfate
absorbs energy in the form of ultra-violet and iodide ions to iodine.

light and changes to atomic oxygen, which

combines with molecular oxygen to form O3. PbS(s) + 4O3(g) PbSO4(s) + 4O2(g)
O2 U.V. light O + O
2KI(aq) + H2O(l) + O3(g) 2KOH(aq) +
O2 + O O3 I2(g) + O2(g)

150

Ozone oxidises nitrogen oxide and gives • The depletion of ozone layer has been most
nitrogen dioxide. pronounced in polar regions, especially
over Antarctica.
NO(g) + O3(g) NO2(g) + O2(g)
• Ozone depletion is a major environmental
Hence the nitrogen oxide emmited from the problem because it increases the amount
of ultraviolet (UV) radiation that reaches
exhaust systems of supersonic jet aeroplanes earth’s surface, thus causing an increase
in rate of skin cancer, eye cataracts and
can bringforth depletion of ozone layer in the genetic as well as immune system damage
among people.
upper atmosphere.
Do you know ?
ii. Bleaching property : Ozone acts as a good Ozone reacts with unsaturated
bleaching agent due to its oxidising nature.
compounds containing double
O3 O + O2 colourless matter bonds to form addition products called
Coloured matter + O ozonides. Ozonides are decomposed by
water or dilute acids to give aldehydes
Ozone bleaches in absence of moisture so it or ketones. This reaction is termed as
is also known as dry bleach. ozonolysis.

iii. Reducing property : Ozone reduces
peroxides to oxides.

e.g. H2O2 + O3 H2O + 2O2
BaO2 + O3 BaO + 2O2

Try this... Internet my friend

a. Ozone is used as bleaching www.britannica.com
agent. Explain. ozone depletion.

b. Why does ozone act as a powerful d. Structure of Ozone : Ozone (O3) is an
oxidising agent ? angular molecule. The two O O bond
lengths in the ozone molecule are identical,
iv. Ozone depletion : Thinning of ozone 128 pm and the O O O bond angle of
layer in upper atmosphere is called ozone about 117°. It is a resonance hybrid of two
depletion. canonical forms.

• The ozone (O3) layer in the upper O O 128 pm O128 pm
atmosphere, absorbs harmful UV OO OO O 117° O
radiations from the sun, thus protecting
people on the earth. e. Uses of ozone :
• Ozone is used for air purification at
• Depletion of ozone layer in the upper
atmosphere is caused by nitrogen oxide crowded places like cinema halls, tunnels,
released from exhausts system of car or railways, etc.
supersonic jet aeroplanes. • In sterilizing drinking water by oxidising
all germs and bacteria.
NO (g) + O3 (g) NO2 (g) + O2 (g) • For bleaching ivory, oils, starch, wax and
delicate fabrics such as silk.
• Depletion (thining) of ozone layer can • In the manufacture of synthetic camphor,
also be caused by chlorofluoro carbons potassium permanganate, etc.

(freons) used in aerosol and refrigerators

and their subsequent escape into the

atmosphere.

151

7.10 Compounds of sulfur : iv. Reaction with Na2SO3 : Sulfur dioxide
reacts with sodium sulfite solution to form
7.10.1 Sulfur dioxide sodium hydrogen sulfite.

a. Preparation : Na2SO3 + H2O(l) + SO2 2NaHSO3

i. From sulfur : Sulfur dioxide gas can be v. Reducing property : Sulfur dioxide acts as
prepared by burning of sulfur in air.
a reducing agent in the presence of moisture.
S(s) + O2 (g) SO2 (g)
• Moist sulfur dioxide reduces ferric salts
ii. From sulfite : In the laboratory sulfur into ferrous salts.

dioxide is prepared by treating sodium sulfite 2Fe3⊕ + SO2 + 2H2O 2Fe2⊕ + SO42 + 4H⊕

with dilute sulfuric acid.

Na2SO3 + H2SO4(aq) Na2SO4+H2O(l)+ SO2(g) • Moist sulfur dioxide decolourises acidified
potassium permangnate (VII) solution.
iii. From sulfides : (Industrial method)
2KMnO4 + 5SO2 + 2H2O K2SO4 +
Sulfur dioxide can be prepared by roasting 2MnSO4 + 2H2SO4
zinc sulfide and iron pyrites.
• Moist sulfur dioxide reduces halogens to
2ZnS(s) + 3O2(g) ∆ 2ZnO(s) + 2SO2(s) halogen acids.

4FeS2(s) + 11O2 (g) ∆ 2Fe2O3(s) + 8SO2(g) I2 + SO2 + 2H2O H2SO4 + 2HI

b. Physical properties of SO2 d. Structure of SO2 : Sulfur dioxide is
angular with O S O bond angle of
i. Sulfur dioxide is a colourless gas with a
119.5°.
pungent smell.

ii. It is poisonous in nature. S S
OO OO
iii. SO2 is highly soluble in water and its
solution in water is called sulfurous acid.

iv. It liquifies at room temperature under a The S O double bond arises from dπ -
pressure of 2 atm and boils at 263 K. pπ bonding. It is a resonance hybrid of two
canonical forms.
c. Chemical Properties :
e. Uses : Sulfur dioxide is used
i. Reaction with Cl2 : Sulfur dioxide reacts
with chlorine in the presence of charcoal • In refining of petroleum and sugar.
(catalyst) to form sulfuryl chloride. • In bleaching wool and silk.
SO2(g) + Cl2(g) charcoal SO2Cl2(l) • As an anti-chlor, disinfectant.
• As a preservative.
ii. Reaction with O2 : Sulfur dioxide is • In the manufacture of H2SO4, NaHSO3.
oxidised by dioxygen in presence of vanadium • Liquid SO2 is used as a solvent to dissolve
(V) oxide to sulfur trioxide.
2SO2(g) + O2(g) V2O5 2SO3(g) a number of organic and inorganic
chemicals.
iii. Reaction with NaOH : Sulfur dioxide
readily reacts with sodium hydroxide solution 7.10.2 Sulfuric acid, H2SO4
to form sodium sulfite.
a. Preparation : Sulfuric acid is
2NaOH + SO2 Na2SO3 + H2O manaufactured by Contact process, which
involves the following three steps.

152

Fig. 7.1 : Flow diagram for manufacture of Sulfuric acid

i. Sulfur or sulfide ore (iron pyrites) on iii. It freezes at 283 K and boils at 611 K.

burning or roasting in air produces sulfur iv. It is highly corrosive and produces severe
dioxide. burns on the skin.

S(s) + O2(g) ∆ SO2 (g) Do you know ?

4FeS2(s) + 11O2(g) ∆ 2Fe2O3(s) + 8SO2(g) Sulfuric acid dissolves in water
with the evolution of a large quantity
ii. Sulfur dioxide is oxidised catalytically of heat. Hence care must be taken while
preparing solution of sulfuric acid from
with oxygen to sulfur trioxide, in the presence concentrated sulfuric acid. Concentrated
H2SO4 must be added slowly to water with
of V2O5 catalyst. 2SO3(g) constant stirring by keeping the beaker in
2SO2(g) + O2 V2O5 water bath.

The reaction is exothermic and reversible

and the forward reaction leads to decrease

in volume. Therefore low temperature (720K)

and high pressure (2 bar) are favourable c. Chemcial Properties :
conditions for maximum yield of SO3.
i. Acidic Property : Sulfuric acid ionises in
iii. Sulfur trioxide gas (from the catalytic aqueous solution in two steps.
converter) is absorbed in concentrated
H2SO4 to produce oleum. H2SO4(aq)+H2O(l) H3O⊕(aq)+HSO4 (aq)
Ka > 10
Dilution of oleum with water gives sulfuric HSO 4(aq)+ H2O(l)
acid of desired concentration. H3O⊕(aq) + SO42 (aq)

SO3(g) + H2SO4 H2S2O7 Ka = 1.2 × 10-2

oleum The greater value of Ka (Ka>10) means that
H2SO4 is largely dissociated into H⊕ and
H2S2O7 + H2O 2H2SO4 HSO4 ions. Thus H2SO4 is a strong acid.

The sulfuric acid obtained by contact process ii. Reaction with metals and nonmetals
(oxidising property) : Metals and nonmetals
is 96 - 98 % pure.

b. Physical properties of H2SO4 : both are oxidised by hot, concentrated sulfuric
i. Sulfuric acid is a colourless, dense, oily
liquid. acid which itself gets reduced to SO2.

ii. It has a density (specific gravity) of 1.84 Cu + 2H2SO4 CuSO4 + SO2 + 2H2O
(Conc.)

g/cm3 at 298 K.

153

S + 2H2SO4 3SO2 + 2H2O d. Uses : Sulfuric acid is a very important
CO2 + 2SO2 + 2H2O industrial chemical. It is used
(Conc.)
• In the manufacture of fertilizers.
C + 2H2SO4 For example, ammonium sulfate,
superphosphate, etc.
(Conc.)
• In the manufacture of pigments, paints and
Remember... dyestuff intermediates.

Oxidizing properties of sulfuric • In petroleum refining.
acid depend on its concentration
and temperature. In dilute solutions, • In detergent industry.
at room temperature, H2SO4 behaves like
HCl, oxidizing metals that stand above • In metallurgy, for cleaning of metals
hydrogen in the e.m.f. series. electroplating and galvanising.

Fe (s) + 2 H⊕ (aq) Fe2⊕ (aq) + H2 (g) • In storage batteries.

Hot, concentrated H2SO4 is a better • As a laboratory reagent.
oxidizing agent than the dilute, cold acid.
• In the manufacture of nitrocellulose
It oxidises metals like copper. products.

iii. Dehydrating property : Concentrated 7.11 Chlorine and compounds of chlorine
sulfuric acid is a strong dehydrating agent.
7.11.1 Chlorine : Chlorine was discovered
Sulfuric acid removes water from sugar and by Scheele, a German Swedish chemist in
carbohydrates. Carbon left behind is called 1774 by the action of HCl on MnO2. In 1810,
sugar charcoal and the process is called Davy established its elementary nature and
charring. suggested the name chlorine on account of its
colour.
C12H22O11 conc. H2SO4 12C + 11H2O (Greek, Chloros = greenish yellow).

iv. Reaction with salts : Concentrated sulfuric a. Preparation :
acid decomposes the salts of more volatile
acids to the corresponding acid 1. Chlorine can be prepared by the oxidation
of hydrochloric acid with any of the following
e.g. NaCl + H2SO4 NaHSO4 + HCl oxidising agents.
KNO3 + H2SO4 KHSO4 + HNO3
CaF2 + H2SO4 CaSO4 + 2HF i. Manganese dioxide :

MnO2 + 4HCl MnCl2 + Cl2 + 2H2O

ii. Potassium permanganate :

Problem 7.10 : What is the action of 2KMnO4 + 16HCl 2KCl + 2MnCl2 +
concentrated H2SO4 on (a) HBr (b) HI 8H2O + 5Cl2

Solution : Concerntrated sulfuric acid 2. Chlorine can also be prepared by the action
oxidises hydrobromic acid to bromine.
of concentrated sulfuric acid on a mixture of

2HBr + H2SO4 Br2 + SO2 + 2H2O sodium chloride (common salt) and manganese

dioxide. The reaction takes place in two steps.

It oxidises hydroiodic acid to iodine. 4NaCl + 4H2SO4 4NaHSO4 + 4HCl

2HI + H2SO4 I2 + SO2 + 2H2O MnO2 + 4HCl MnCl2 + 2H2O + Cl2

4NaCl + MnO2 + 4H2SO4 4NaHSO4 +
MnCl2 + 2H2O + Cl2

154

b. Manufacture of chlorine : iv. Reaction with NH3 : Chlorine when reacted
with excess of ammonia gives ammonium
i. Deacon’s process : chloride and nitrogen.

Chlorine is manufactured by the oxidation of 8NH3 + 3Cl2 6NH4Cl + N2
hydrogen chloride gas by atmospheric oxygen
in the presence of CuCl2 as catalyst at 723 K. (Excess)
4HCl + O2 CuCl2 2Cl2 + 2H2O
Excess of chlorine reacts with ammonia to
ii. Electrolytic process : give nitrogen trichloride (explosive).

By the electrolysis of brine (concentrated NaCl NH3 + 3Cl2 NCl3 + 3HCl
solution), chlorine is liberated at the anode. (Excess)

NaCl Na⊕ + Cl v. Reaction with alkali : Chlorine reacts with
cold and dilute alkali to produce a mixture of
Cathode: 2H2O + 2e H2 + 2OH chloride and hypochlorite. When reacted with
Na⊕ + OH NaOH hot concentrated alkali, chloride and chlorate
are produced.

Anode : Cl Cl + e 2NaOH + Cl2 NaCl + NaOCl + H2O

Cl + Cl Cl2 (Cold and dilute)

c. Physical Properties of Chlorine : 6NaOH + 3Cl2 5NaCl + NaClO3 + 3H2O

i. Chlorine is a greenish-yellow gas having (Hot and conc.)
pungent and suffocating odour.
Chlorine when reacted with dry slaked lime
ii. It is poisonous in nature. gives bleaching powder.

iii. It can be easily liquified into a greenish 2Ca(OH)2 + 2Cl2 Ca(OCl)2 + CaCl2 +
yellow liquid, which boils at 293 K. 2H2O

iv. It dissolves in water to give chlorine water. vi. Reaction with hydrocarbons : Chlorine
reacts with saturated hydrocarbons to give
v. It is 2-5 times heavier than air. substitution products and with unsaturated
hydrocarbons gives addition products.
d. Chemical properties of chlorine :

i. Reaction with metals : Chlorine reacts with CH4 + Cl2 U.V CH3Cl + HCl
metals to form chlorides. Methane Methylchloride

2Al + 3Cl2 2AlCl3 H2C = H2C + Cl2 CH2-CH2
Ethene Cl Cl
2Na + Cl2 2NaCl 1,2 - Dichloroethane

2Fe + 3Cl2 2FeCl3 vii. Oxidising property : Chlorine oxidises
ferrous salts to ferrric salts and sulfites to
ii. Reaction with nonmetals : Chlorine reacts sulfates.

with nonmetals to form their chlorides.

P4 + 6Cl2 4PCl3 2FeSO4 + H2SO4 + Cl2 Fe2(SO4)3 + 2HCl

S8 + 4Cl2 4S2Cl2 Na2SO3 + Cl2 + H2O Na2SO4 + 2HCl

iii. Affinity for hydrogen : Chlorine has great It oxidises sulfur dioxide to sulfur trioxide

affinity for hydrogen. It reacts with hydrogen and iodine to iodate. In presence of water they

and compounds containing hydrogen to form form sulfuric acid and iodic acid respectively.

HCl. 2HCl SO2 + 2H2O + Cl2 H2SO4 + 2HCl
2HCl + S I2 + 6H2O + 5Cl2 2HIO3 + 10HCl
H2 + Cl2
H2S + Cl2

155

viii. Bleaching Property : Chlorine requires 7.11.2 Hydrogen Chloride : Hydrogen
the presence of moisture (water) for bleaching. chloride was prepared by Glauber in 1648
It liberates nascent oxygen from water which by heating common salt with concentrated
is responsible for its oxidising and bleaching sulfuric acid. Davy in 1810 showed that it is
property a compound of hydrogen and chlorine.

Cl2 + H2O HCl + HOCl a. Preparation : In the laboratory, hydrogen
chloride is prepared by heating sodium
HOCl HCl + [O] chloride (common salt) with concentrated
sulfuric acid.
Chlorine bleaches vegetable matter or coloured
organic matter in the presence of moisture to NaCl + H2SO4 420 K NaHSO4 + HCl
colourless matter. NaHSO4 + NaCl 420 K Na2SO4 + HCl

Coloured organic matter + [O] Colourless HCl gas can be dried by passing it
organic matter through concentrated sulfuric acid.

e. Uses :

Chlorine is used a. Physical properties of HCl

• For purification (sterilizing) of drinking i. Hydrogen chloride is a colourless and
water. pungent smelling gas.

• For bleaching wood pulp required for ii. It can be easily liquified to a colourless
manufacture of paper and rayon, bleaching liquid (b.p. 189 K) which freezes to a white
cotton and textiles. crystalline solid (m.p. 159 K)

• For extraction of metals like gold and iii. It is highly soluble in water.
platinum.
Chemical properties :
• In the manufacture of dyes, drugs and i. Acidic property : Hydrogen chloride is
organic compounds such as CCl4, CHCl3, highly soluble in water and ionises as follows
DDT, refrigerants, etc. :

• In the preparation of poisonous gases such HCl (g) + H2O (l) H3O⊕ (aq) + Cl (aq)
as phosgene (COCl2), tear gas (CCl3NO2),
mustard gas (ClCH2CH2SCH2CH2Cl). Ka = 107

Try this... The aqueous solution of HCl gas is called

i. Give the reasons for bleaching hydrochloric acid. High value of dissociation
action of chlorine.
ii. Name the two gases used in war. constant (Ka) indicates that it is a strong acid
in water.

Do you know ? ii. Reaction with NH3 : Hydrochloric acid
reacts with ammonia and gives white fumes
Bleachingbychlorineis permanent. of ammonium chloride.
It bleaches cotton fabrics, wood
pulp, litmus, etc. However chlorine is not NH3 + HCl NH4Cl
used to bleach delicate materials such as
silk, wool etc. as it is a strong bleaching iii. Reaction with noble metals : When three
and oxidising agent. This dual action will parts of concentrated HCl and one part of
damage the base material. concentrated HNO3 are mixed, aqua regia is
formed.

Noble metals like gold, platinum get dissolved
in aqua regia.

156

Au + 4 H⊕ + NO3 + 4Cl AuCl4 + NO neutral molecules are called Interhalogen
3Pt + 16 H⊕ + 4NO3 + 18Cl + 2H2O compounds. For example, CIF, BrF3.

3PtCl62 + An interhalogen compound is a compound
4NO + 8H2O formed by combination of atoms of different
halogens. The interhalogen compound
Can you recall ? is regarded as the halide of the more
electropositive halogen. A given halogen
• Which type of bonds do forms an interhalogen compound only with
halogens form with other the halogen having lesser electronegativity.
elements?
Do you know ?
• Does BrF5 obey the octet rule?
• What is the oxidation state of Br in The ions formed by combination
of different halogens are called
BrF5? polyhalide ions or interhalogen ions, For
• How many electrons do halogens example, K⊕[Cl3 ], [NH4]⊕ [I5] , which
contain Cl3 and I5 ions.
require to complete their octet ?

• What is the shape of CIF3?

7.12 Interhalogen compounds Use your brain power

We have seen that all halogen molecules • Chlorine and fluorine combine
are diatomic. They form binary compounds to form interhalogen compounds.
with hydrogen, with oxygen they form The halide ion will be of chlorine or
oxyacids, and with halogens they form fluorine ?
interhalogen compounds.
• Why does fluorine combine with other
Although all halogens belong to the same halogens to form maximum number
group they have different electronegativites. of fluorides ?
Due to this difference in electronegativity
two or more halogen atoms combine to form
species which may be ionic or neutral. The

Table 7.9 Various types of interhalogen compounds

Element Fluoride Chloride Bromide Iodide

Chlorine CIF, CIF3, CIF5 - - -

Bromine BrF, BrF3, BrF5 BrCl - -
Iodine IF, IF3, IF5, IF7 ICl, ICl3 IBr -

Do you know ? 7.12.1 Classification : Depending on their
composition, interhalogen compounds are
ICl Iodine monochloride in glacial classified into four types.
acetic acid called Wijs solution is
used in determination of iodine value of Type Example
an oil. XX' CIF, BrF, BrCl, ICl, IBr
XX'3 CIF3, BrF3, IF3
XX'5 CIF5, BrF5,IF5
XX'7 IF7

157

In the general formula XX'n. X is Use your brain power
the halogen having larger size and is more
electropositive. X' is the halogen having • Which halogen has tendency
smaller size and more electronegativity. to form more interhalogen
compounds?
7.12.1 General characteristics of

interhalogen compounds Use your brain power
• Which will be more reactive ?
1. The compound is considered as the
halide of X. For example, ClF. Here the a. ClF3 or ClF, b. BrF5 or BrF
halogen having larger size is chlorine, it is
more electropositive than F and hence the • Complete the table

interhalogen compound is named as chlorine Formula Name
monofluoride. (n) is the number of atoms of ClF Chlorine monofluoride
X' attached to X. ClF3
Chlorine penta fluoride
As the ratio [radius of X : radius of X'] BrF
increases the value of n also increases. Bromine penta fluoride
2. Interhalogen compounds have even number ICl
of atoms 2, 4, 6, 8. For example, ClF3 has 4 ICl3
atoms.
3. The properties of interhalogen compounds Table 7.10 : States of Interhalogen compounds
are generally intermediate between those of at 25°C
the halogens from which they are made. XX'
4. The central halogen exhibits different
oxidation states in different interhalogen ClF Colorless gas
compounds.

5. Number of X' atoms in the compounds is BrF Pale brown gas
always odd. BrCl
6. They are all diamagnetic . ICl Gas

Ruby red solid (α- form)
Brown red soid (β - form)

Use your brain power IBr Black solid
• What will be the names of XX3'
ClF3
the following compounds : BrF3 Colorless gas
ICl, BrF. IF3
• Which halogen (X) will have maximum ICl3 Yellow green liquid
number of other halogen (X') attached?
IF5 Yellow powder
Do you know ? BrF5
XX' compounds are more reactive ClF5 Orange solid dimerises
than X2 or X'2. In the X-X' bond X' to form (I2Cl6 having Cl -
is more electronegative than X, while IF7 bridges)
in X2 and X'2 both atoms have same
electronegativity, hence the X-X' bond XX5'
energy is less than the X-X or X'-X' bond Colorless gas at R. T. but
energy. solid below 77 K
Colorless liquid

Colorless liquid

XX7'
Colorless gas

158

Methods of Preparation of Interhalogen compounds

Method XX’ XX’3
1.Direct combination Both in equal volumes
Cl2+3F2 (excess) 2ClF3
Cl2+F2 523 K 2ClF Br2+3F2 2BrF3
Cu-vessel
I2+3Cl2(excess) 2ICl3

I2+ Cl2 HCl + HNO3 2ICl
boil

2.Reaction of halogen with Br2+Cl2 2BrCl Br2+ ClF3 2BrF3+ BrCl
interhalogen compounds. Br2+ BrF3 3BrF

Special reaction for ICl I2+KClO3 ∆ ICl+KIO3 -------

Use your brain power Do you know ?

In the above special reaction ICl in liquid state, undergoes
for ICl, identify the oxidant and autoionization like water, to form
reductant ? Denote oxidation states of a cation and an anion.
the species.
2 ICl I⊕ + ICl2

(solvent cation) (solvent anion)

Some Properties of Interhalogen compounds

Reaction/Property XX’ XX’3

1 Thermal stability ClF > ICl > IBr > BrCl > BrF

2 Hydrolysis Gives BrCl+H2O HOBr+HCl ICl3+H2O
oxoacids 5ICl+3H2O HIO3+5HCl+2I2 HIO3+5HCl+ICl

3 Disproportionation/ BrF 2Br2 + BrF5 2ClF3 ClF2⊕+ClF4
Autoionisation ICl3 341 K ICl + Cl2

4 Flourination --------- U+ClF3 UF6(l)+ClF(g)

5 Addition across I Cl

olefins H2C = CH2 + ICl H C C H

HH

Uses of interhalogen compound

XX' XX'3
For Preparation of polyhalides
ICl is used to determine iodine value of
oils

As catalyst for oxidation of As(III) As fluorinating agent

For Preparation of polyhalides As nonaqueous solvent

159

Table 7.11 Structures of some Interhalogen Compounds

Formula Name Structure Shape
ICl
ClF3 Iodine monochloride Cl I Cl Linear
BrF3
Chloride trifluoride F Cl F Bent T- shaped
BrF5 Bromine trifluoride Bent T- shaped
89°
ClF5
F

F
Br 86F°

F

Bromine pentafluoride FF Square Pyramidal
Br

FF
F

Chlorine pentafluoride FF Square pyramidal
Cl

FF
F

Table 7.12 Oxidation states of central halogen We have studied earlier in this chapter
atom in interhalogen compounds. that group 18 elements have very high
ionisation energies and due to this property
O.S. No. of lone Examples they are unreactive. Each noble gas atom
central pairs of has a completely filled valence electron shell
Halogen electrons which makes it inert.

+7 0 IF7 The first ionization potential decreases
down the group, hence heavier noble gases
+5 1 ClF5, BrF5, Kr, Xe and Rn can form compounds due to
IF5 low ionization energy.

+3 2 ClF3, BrF3,
IF3, I2Cl6

ClF, BrF, IF, Do you know ?

+1 3 BrCl, ICl, First true compound of noble
gas was made in 1962 by Neil
IBr Bartelt and Lohman.

7.13 Compounds of Xenon Xe + 2[PtF6] [XeF]⊕ [Pt2F11]

Can you recall ? Only Xenon reacts directly with
• What is the correlation
fluorine to form Xenon fluorides.
between ionization energies
and reactivity of elements? Remember...
• Trends in ionization energy down a XeF2, XeF4, XeF6 are stable
group.
fluoride of xenon.

Xenon also forms compounds with
oxygen, such as XeO3, XeOF2, XeOF4,
XeO2F2.

160

7.13.1 Xenon fluorides ii. Reaction with PF5 : XeF2 forms adducts
on reaction with PF5.
a. Preparation of Xenon Fluorides : Xenon
fluorides are generally prepared by direct XeF2 + PF5 XeF2.PF5
reaction of xenon and fluorine in different
ratios and conditions, such as temperature, 7.13.2 Xenon trioxide
electric discharge and photochemical reaction.
a. Preparation of Xenon trioxide (XeO3) :
i. Xe + F2 sealed XeF2
Ni tube, 4000C Fluorides of Xenon react with water to
form XeO3.
Ni tube, 4000C
iv. Xe + F2 5-6 atmos XeF4 3XeF4 + 6H2O 2Xe + XeO3 + 12 HF
1 : 5 +
1 1 O2
v. Xe + 2F2 electric discharge XeF4 2
-80 0C
XeF6 + 3H2O XeO3 + 6HF
vi. Xe + 3F2 electric discharge XeF6
low temp. 7.13.3 Oxyfluorides of Xenon : Xenon forms

b. Important chemical reactions of XeF2 the following oxyfluorides :
i. Hydrolysis : XeF2 undergoes hydrolysis
XeOF2, XeO2F2, XeO3F2, XeO2F4
to form HF

2XeF2 + 2H2O 4HF + 2Xe + O2

Table 7.13 : Structure of Xenon Compounds

Sr. No. Formula Oxidation state of Xe Structure

1. XeF2 +2 Linear

2. XeF4 +4 Square planar
3. XeF6 +6 Distorted octahedral

4. XeO3 +6 Trigonal pyramidal
5. XeOF4 +6 Square pyamidal

161

a. Preparation

i. Partial hydrolysis of Xenon fluorides Use your brain power
yields different oxyfluorides.
What are the missing entries ?

XeF4 + H2O 80 0C XeOF2 + 2HF Formula Name

XeF6 + H2O XeOF4 + 2HF XeOF2 Xenon mono
oxyfluoride
ii. Reaction of Xenon oxyfluoride with SiO2
or hydrolysis yields xenon dioxydifluoride. Xenon dioxydifluoride

2XeOF4 + SiO2 2XeO2F2 + SiF4 XeO3F2
XeOF4 + H2O XeO2F2 + 2HF
XeO2F4

Element Table 7.14 : Uses of helium, neon and argon
Helium
Uses
Neon i. Mixture of He and O2 is used for artificial breathing of asthma patients.
Argon ii. Mixture of He and O2 is used for respiration by sea divers.
iii. For filling balloons, a mixture of helium (85%) and hydrogen (15 %) is

used.

iv. Helium is used for producing inert atmosphere required for welding
purpose and metallury of some metals.

v. Liquid helium is used for producing low temperature required for research.

vi. In low temperature gas thermometry, for production of lasers.

vii. Used to pressurise fuel tanks of liquid fueled rockets.

viii. Used as shielding gas for arc welding.

ix. In supersonic wind tunnels.

x. Helium nucleus is used as a bombarding particle for disintegration of
atoms.

xi. Used for magnetic resonance imaging.

i. In Neon discharge lamps and signs. These signs are visible from the long
distances and also in mist or fog.

ii. Mixture of Ne and He is used in certain protective electrical devices such
as voltage stabilizers and current rectifiers.

iii. For production of lasers.

iv. In fluorescent tubes.
i. For producing inert atmosphere in welding and steel production.

ii. Mixture of 85 % Ar and 15 % N2 is filled in electric bulb to increase life
of filament.

iii. In filling fluorescent tubes and radio valves.

iv. It is mixed with neon to get lights of various colors.

162

Exercises

1. Select appropriate answers for the x. Chlorine reacts with excess of fluorine to
following.
form.

i. Which of the following has highest A. ClF B. ClF3
C. ClF2 D. Cl2F3
electron gain enthalpy ?

A. Fluorine B. Chlorine xi. In interhalogen compounds, which of the

C. Bromine D. Iodine following halogens is never the central

ii. Hydrides of group 16 are weakly acidic. atom.
The correct order of acidity is
A. I B. Cl C. Br D. F
A. H2O > H2S > H2Se > H2Te
B. H2Te > H2O > H2S > H2Se xii. Which of the following has one lone pair
C. H2Te > H2Se > H2S > H2O
D. H2Te > H2Se > H2O > H2S of electrons ?

iii. Which of the following element does not A. IF3 B. ICl
show oxidation state of +4 ? C. IF5 D. ClF3

A. O B. S C. Se D. Te xiii. In which of the following pairs, molecules

are paired with their correct shapes ?

iv. HI acid when heated with conc. H2SO4 A. [I3] : bent
forms B. BrF5 : trigonal bipyramid
C. ClF3 : trigonal planar
A. HIO3 B. KIO3 D. [BrF4] : square planar
C. I2 D. KI
xiv. Among the known interhalogen

v. Ozone layer is depleted by compounds, the maximum number of

A. NO B. NO2 atoms is
C. NO3 D. N2O5
A. 3 B. 6 C. 7 D. 8

vi. Which of the following occurs in liquid 2. Answer the following.

state at room temperature ? i. Write the order of the thermal stability of
the hydrides of group 16 elements.
A. HIO3 B. HBr
C. HCl D. HF ii. What is the oxidation state of Te in TeO3 ?
iii. Name two gases which deplete ozone
vii. In pyrosulfurous acid oxidation state of
sulfur is layer.

A. Only +2 B. Only +4 iv. Give two uses of ClO2
C. +2 and +6 D. Only +6 v. What is the action of bromine on

viii. Stability of interhalogen compounds magnesium metal ?
follows the order
vi. Write the names of allotropic forms of
A. BrF > IBr > ICl > ClF > BrCl selenium.
B. IBr > BeF > ICl > ClF > BrCl
C. ClF > ICl > IBr > BrCl > BrF vii. What is the oxidation state of S in H2SO4.
D. ICl > ClF > BrCl > IBr > BrF viii. The pKa values of HCl is -7.0 and that of

ix. BrCl reacts with water to form HI is -10.0. Which is the stronger acid ?

A. HBr B. Br2 + Cl2 ix. Give one example showing reducing
C. HOBr D. HOBr + HCl property of ozone.

163

x. Write the reaction of conc. H2SO4 with viii. What is the action of chlorine on the
sugar. following

xi. Give two uses of chlorine. a. Fe b. Excess of NH3

xii. Complete the following. ix. How is hydrogen chloride prepared from
sodium chloride ?
1. ICl3 + H2O ........ + ...... + ICl

2. I2 + KClO3 ....... + KIO2 x. Draw structures of XeF6, XeO3, XeOF4,
3. BrCl + H2O ....... + HCl XeF2.

4. Cl2 + ClF3 ........ xi. What are inter-halogen compounds ? Give
two examples.
5. H2C = CH2 + ICl .......

6. XeF4 + SiO2 ....... + SiF4 xii. What is the action of hydrochloric acid on
7. XeF6 + 6H2O ........ + HF
the following ?
8. XeOF4 + H2O ....... + HF
a. NH3 b. Na2CO3
xiii. Match the following
xiii. Give two uses of HCl.
A B

XeOF2 Xenon trioxydifluoride xiv. Write the names and structural formulae
XeO2F2 Xenon monooxydifluoride of oxoacids of chlorine.
XeO3F2 Xenon dioxytetrafluoride
XeO2F4 Xenon dioxydifluoride xv. What happens when
a. Cl2 reacts with F2 in equal volume at
xiv. What is the oxidation state of xenon in the
437 K.
following compounds. b. Br2 reacts with excess of F2.

XeOF4, XeO3, XeF6, XeF4, XeF2. xvi. How are xenon fluorides XeF2, XeF4 and
XeF6 obtained ? Give suitable reactions.
3. Answer the following.
xvii. How are XeO3 and XeOF4 prepared ?
i. The first ionisation enthalpies of S, Cl
and Ar are 1000, 1256 and 1520 kJ/mol-1, xviii. Give two uses of neon and argon.
respectively. Explain the observed trend.
xix. Describe the structure of Ozone. Give two
ii. “Acidic character of hydrides of group uses of ozone.
16 elements increases from H2O to H2Te”
Explain. xx. Explain the trend in following atomic
properties of group 16 elements.
iii. How is dioxygen prepared in laboratory
from KClO3 ? i. Atomic radii ii. Ionisation enthalpy

iv. What happens when iii. Electronegativity.
a. Lead sulfide reacts with ozone (O3).
b. Nitric oxide reacts with ozone. 4. Answer the following.

v. Give two chemical reactions to explain i. Distinguish between rhombic sulfur and
oxidizing property of concentrated H2SO4. monoclinic sulfur.

vi. Discuss the structure of sulfure dioxide. ii. Give two reactions showing oxidising
property of concentrated H2SO4.
vii. Fluorine shows only -1 oxidation state
while other halogens show -1, +1, +3, +5 iii. How is SO2 prepared in laboratory
and +7 oxidation states. Explain. from sodium sulfite? Give two physical
properties of SO2.

iv. Describe the manufacturing of H2SO4 by
contact process.

164

8. TRANSITION AND INNER TRANSITION ELEMENTS

Do you know ? 8.2 Position in the periodic table

In which block of the modern The transition elements are placed in
perodic table are the transition and the periods 4 to 7 and groups 3 to 12 those
inner transition elements placed ? constitute 3d, 4d, 5d and 6d series (Fig.8.1).

8.1 Introduction They are placed at the centre with

The transition elements belong to s block on one side and p on the other. The

d block of the periodic table. As per IUPAC electropositivity, reactivity and other properties

convention the transition metal atom has an show a gradual change from s block to p block

incomplete d-subshell or it give cations with through those of the d block elements.

incomplete d subshell. They exhibit properties 8.3 Electronic configuration
between those of s and p block elements. The
transition elements of the modern periodic The 3d series begins with Sc (Z = 21) and
table appear as groups 3 to 12 or as four long ends with Zn (Z=30). Argon, Ar is the noble
periods. The (n-1) d-orbital is comprised of gas preceding to 3d series and its electronic
successively filled in each element, where ‘n’ configuration is 1s2 2s2 2p6 3s2 3p6. Calcium
is the ultimate or valence shell. The 3d series (Z = 20) belonging to ‘s’ block of 4th period
is comprised of elements from scandium has electronic configuration 1s2 2s2 2p6 3s2 3p6
(Z=21) to zinc (Z=30), 4d series has elements 4s2. Hence 21st electron in scandium (Z = 21)
from yttrium (Z=39) to cadmium (Z=48), 5d enters in the available 3d orbital. Electronic
series from lanthanum (Z=57) to mercury configuration of Sc is written as 1s2 2s2 2p6 3s2
(Z=80) without those from cerium to lutecium, 3p6 3d1 4s2 or also can be represented as [Ar]
and 6d series has actinium to curium without 3d1 4s2.

those fromthorium to lawrentium. The general

electronic configuration of transition elements

is (n-1)d1-10 ns1-2.

Table 8.1 Shows the four transition series elements

Group 3 4 5 6 7 8 9 10 11 12
d series

3d Sc(21) Ti(22) V(23) Cr(24) Mn(25) Fe(26) Co(27) Ni(28) Cu(29) Zn(30)

4d Y(39) Zr(40) Nb(41) Mo(42) Tc(43) Ru(44) Rh(45) Pd(46) Ag(47) Cd(48)

5d La(57) Hf(72) Ta(73) W(74) Re(75) Os(76) Ir(77) Pt(78) Au(79) Hg(80)

6d Ac(89) Rf(104) Db(105) Sg(106) Bh(107) Hs(108) Mt(109) Ds(110) Rg(111) Cn(112)

165

s - block p - block

1 18
1
2 13 14 15 16 17

2 d - block
5 6 7 8 9 10 11 12
3 34
4
5

6

7

Series
and
f - block

Fig. 8.1 Position of d-block elements in the modern periodic table

The electronic configuration of the 8.3.1 Electronic configuration of chromium
elements of 3d series is given in Table 8.2. and copper

Since zinc has completely filled Table 8.2 indicates that the expected
(n - 1)d orbital in the ground state (3d10,4s2) electronic configuration of chromium (Z = 24)
and (3d10) in its common oxidation state +2, differs from the observed configuration.
it is not regarded as transition element. On the
same ground, cadmium and mercury from 4d This can be explained on the basis of
and 5d series are not considered as transition the concept of additional stability associated
elements. Copper in the elementary state (3d10 with the completely filled and half filled
4s1) contains filled 3d orbitals but in the +2 subshells.
oxidation state it has partly filled 3d orbital
(3d9), hence copper is a transition element. Remember...

General electronic configuration of Any subshell having a half filled
four series of d-block elements of periodic or completely filled electronic
table can be represented as given below: configuration has extra stability.

i. 3d series : [Ar] 3d1-10 4s2 The general electronic configuration
of the elements of the 3d series is 3d1-10 4s2
ii. 4d series : [Kr] 4d1-10 5s0-2 with the exceptions of Cr and Cu. The 3d and
4s orbitals are close in energy and in order to
iii. 5d series : [Xe] 5d1-10 6s2 gain extra stability the last electron instead of
occupying 4s orbital occupies the 3d orbital
iv. 6d series : [Rn] 6d1-10 7s2 that assigns Cr the 3d5, 4s1 and Cu 3d10, 4s1
configuration.

166

Table 8.2 Electronic configuration of 3d series of d-block elements

Elements Symbol At. No. Expected electronic Observed electronic
configuration configuration
Scandium Sc 21 [Ar] 3d1 4s2 [Ar] 3d1 4s2
Titanium Ti 22 [Ar] 3d2 4s2 [Ar] 3d2 4s2
Vanadium V 23 [Ar] 3d3 4s2 [Ar] 3d3 4s2
Chromiun Cr 24 [Ar] 3d4 4s2 [Ar] 3d5 4s1
Manganese Mn 25 [Ar] 3d5 4s2 [Ar] 3d5 4s2
Fe 26 [Ar] 3d6 4s2 [Ar] 3d6 4s2
Iron Co 27 [Ar] 3d7 4s2 [Ar] 3d7 4s2
Cobalt Ni 28 [Ar] 3d8 4s2 [Ar] 3d8 4s2
Nickel Cu 29 [Ar] 3d9 4s2 [Ar] 3d10 4s1
Copper Zn 30 [Ar] 3d10 4s2 [Ar] 3d10 4s2
Zinc

Use your brain power
Fill in the blanks with correct outer electronic configurations.

Z Y Zr Nb 2nd series Ru Rh Pd Ag Cd
valence 39 40 41 44 45 46 47 48
4d15s2 4d25s2 4d45s1 Mo Tc 4d85s1 4d105s0
shell 42 43 Os 4d105s2
electronic La Hf Ta 76 Ir Pt
configuration 57 72 73 4d65s1 5d66s2 77 78 Au Hg
5d16s2 5d36s2 5d76s2 5d96s1 79 80
Z Rf 3rd series Hs
valence shell Ac 104 Db 108 Mt Ds 5d106s2
89 6d27s2 105 W Re 6d67s2 109 110
electronic 6d17s2 74 75 6d77s2 6d87s2 Rg Uvb
configuration 5d46s2 111 112

Z 4th series 6d107s2
valence shell
Sg Bh
electronic 106 107
configuration 6d47s2 6d57s2

Table 8.3: Oxidation states of first transition series elements

Elements Outer Oxidation states Elements Outer Oxidation states
elecronic electronic
Sc configuration +2, +3 Fe configuration +2, +3, +4, +5, +6
Ti +2, +3, +4 Co +2, +3, +4, +5
V 3d1 4s2 +2, +3, +4, +5 Ni 3d6 4s2 +2, +3, +4
Cr +2, +3, +4, +5, +6 Cu +1, +2
Mn 3d2 4s2 +2, +3, +4, +5, +6, +7 Zn 3d7 4s2 +2

3d3 4s2 3d10 4s2

3d5 4s1 3d10 4s1

3d5 4s2 3d10 4s2

167

Remember... From Table 8.3 it is clear that as the number
Electronic configuration of Cr of unpaired electrons in 3d orbitals increases,
the number of oxidation states shown by the
is [Ar] 3d5 4s1 and element also increases. Scandium has only one
Cu is : [Ar]3d10, 4s1. unpaird electron. It shows two oxidation states
while manganese with 5 unpaired d electrons
8.4 Oxidation states of first transition shows six different oxidation states.
series
The elements which give the greatest
One of the notable features of transition number of oxidation states occur in or near the
elements is the great variety of oxidation states middle of the series. Manganese, for example,
they show in their compounds. Table 8.3 lists shows oxidation states from +2 to +7.
the common oxidation states of the first row 8.5 Physical properties of first transition
transition elements. series : All transition elements are metals
and show properties that are characteristic of
Can you tell ? metals. They are hard, lustrous, malleable,
Which of the first transition series ductile and form alloys with other metals. They
elements shows the maximum are good conductors of heat and electricity.
number of oxidation states and why? Except Zn, Cd, Hg and Mn, all the other
transition elements have one or more typical
Which elements in the 4d and 5d series will metallic structures at ambient temperature.
show maximum number of oxidation states? These transition metals (with the exception
of Zn, Cd and Hg) are very hard and have
Loss of 4s and 3d electrons progressively low volatility. They possess high melting and
leads to formation of ions. The transition boiling points.
elements display a variety of oxidation states
in their compounds. Loss of one 4s electron Fig. 8.2 : Trends in melting points of transition
leads to the formation of M⊕ ion, loss of two elements
4s electrons gives a M2⊕ ion while loss of
unpaired 3d and 4s electrons gives M3⊕, M4⊕
ions and so on. Some examples are as shown
in Table 8.4

Try this...

Write the electronic configuration
of Mn6⊕, Mn4⊕, Fe4⊕, Co5⊕, Ni2⊕

Elements Table 8.4: Electronic configuration of various ions of 3d elements Cu Zn
Atomic no : Sc Ti V Cr Mn Fe Co Ni 29 30
21 22 23 24 25 26 27 28
Species Valence shell Electronic Configuration
M
M⊕ 3d1 4s2 3d2 4s2 3d3 4s2 3d5 4s2 3d5 4s2 3d6 4s2 3d1 4s2 3d8 4s2 3d10 4s1 3d10 4s2
M2⊕ 3d1 4s1 3d2 4s1 3d3 4s1 3d5 3d5 4s1 3d6 4s1 3d7 4s1 3d8 4s1 3d10 4s0 3d10 4s1
M3⊕ 3d4
3d2 3d2 3d3 3d3 3d5 3d6 3d7 3d8 3d9 3d10
[Ar] 3d1 3d2 3d4 3d5 3d6 3d7 - -

168

8.6 Trends in atomic properties of the first atomic radii in nm
transition series
Fig. 8.3 : Trends in atomic radii of d block
8.6.1 Atomic and ionic radii elements

Atomic radii of the elements of the Ionic radii of transition elements show
transition series decrease gradually from left the same trend as of the atomic radii (Table
to right (Fig. 8.3 and Table 8.5). As we move 8.5)
across a transition series from left to right the
nuclear charge increases by one unit at a time. The elements of first transition series
The last filled electron enters a penultimate show variable oxidation states. The trends
(n-1)d subshell. However, d orbitals in an in ionic radii, thus, can be studied with (i)
atom are less penetrating or more diffused and, elements having same oxidation state or (ii)
therefore d electrons offer smaller screening considering various oxidation states of the
effect. The result is that effective nuclear same element.
charge also increases as the atomi number
increases along a transition series. Hence (i) For the same oxidation state, with an
the atomic radii decrease gradually across a increase of nuclear charge a gradual decrease
transition series from left to right. in ionic radii was observed. The trend is
pronounced for the divalent ions of the first
The explanation for the minor variation transition series (Cr2⊕ - 82 pm, Cu2⊕ - 72 pm).
in atomic radii within a particular transition
series is out of the scope of this textbook.

Table 8.5 Atomic properties of first transition series elements
Element
(M) Atomic Density Atomic/ionic radius (pm) Ionisation

Sc number (g/cm3) M M2⊕ M3⊕ enthalpy (kJ/mol)
Ti (Z)
V
Cr 21 3.43 164 - 73 631
Mn
Fe 22 4.1 147 - 67 656
Co
Ni 23 6.07 135 79 64 650
Cu
Zn 24 7.19 129 82 62 653

25 7.21 127 82 65 717

26 7.8 126 77 65 762

27 8.7 125 74 61 758

28 8.9 125 70 60 736

29 8.9 128 73 - 745

30 7.1 137 75 - 906

169

(ii) The oxidation states of the same element The atoms of elements of third
shows difference of one unit such as M⊕, M2⊕, transition series possess filled 4f- orbitals. 4f
M3⊕, M4⊕ and so on. With higher oxidation state orbitals show poor shielding effect on account
the effective nuclear charge also increases and of their peculiar diffused shape. As a result,
hence, decrease of ionic radii can be observed the valence electrons experience greater
from M2⊕ to M3⊕ (Table 8.5). Ionic radii of nuclear attraction. A greater amount of energy
transition elements are smaller than ionic radii is required to ionize elements of the third
of representative elements of the same period. transition series. The ionisation enthalpies of
the elements of the third transition series are,
8.6.2 Ionisation Enthalpy : The ionisation therefore much higher than the first and second
enthalpies of transition elements are series (Fig.8.4).
intermediate between those of s-block or
p-block elements. This suggests that transition 8.6.3 Metallic character : Low ionization
elements are less electropositive than elements enthalpies and vacant d orbitals in the
of group 1 and 2. Depending on the conditions, outermost shell are responsible for the metallic
they form ionic or covalent bonds. Generally in character of the transition elements. These
the lower oxidation states these elements form favour the formation of metallic bonds and
ionic compounds while in the higher oxidation thus these elements show typical metallic
states they form covalent compounds. properties. The hard nature of these elements
suggests the formation of covalent bonds in
Ionisation enthalpies shown in them. This is possible due to the presence of
Table 8.6 reveal that for a given element there unpaired (n-1)d electrons in these elements.
is substantial increase from the first ionisation
enthalpy IE1 to the third ionisation enthalpy Nearly all transition metals have simple
IE3. hexagonal closed packed (hcp), cubic closed
packed (ccp) or body centered cubic (bcc)
As we move across the transition series, lattices which are characteristic of true
slight variation is observed in the successive metals (You have learnt more about this in
enthalpies IE1, IE2, IE3 of these elements (Table Chapter 1).
8.6).

Remember...
Hardness, high melting points

and metallic properties of the
transition elements indicate that the
metal atoms are held strongly by metallic
bonds with covalent character.

Fig. 8. 4: Trends in first ionisation enthalpies of
d block elements

Table 8.6 Ionisation enthalpies of first transition series elements

Element Sc Ti V Cr Mn Fe Co Ni Cu Zn
IE
IE1 632 659 650 652 717 762 756 736 744 906
IE2 1245 1320 1376 1635 1513 1563 1647 1756 1961 1736
IE3 2450 2721 2873 2994 3258 2963 3237 3400 3560 3838

(IE = Ionisation Enthalpy in kJ/mol)

170

In all the transition series the melting Each unpaired electron gives rise to a
points steadily increase upto d5 configuration. small magnetic field (magnetic moment) due
Cr, Mo and W show highest melting points to its spin angular momentum and orbital
in their respective series. Mn and Tc display angular momentum. In case of the first row
anomalous values of melting points. After this transition elements, the contribution from
with increasing atomic number the melting the orbital angular momentum is quenched
point decreases regularly. and hence, can be neglected. The spin-only
formula for magnetic moment is :
8.6.4 Magnetic Properties :
m = n(n + 2) BM
Can you recall ?
where n is the number of unpaired
1. What happens when magnetic electrons and m is the magnetic moment
field is applied to substances ? expressed in Bohr Magneton (BM). A single
unpaired electron has magnetic moment
2. What is meant by the terms m = 1.73 BM.
paramagnetism and diamagnetism ?
From the magnetic moment (m)
The compounds of transition metals measurements of the metal complexes of the
exhibit magnetic properties due to the unpaired first row transition elements, the number of
electrons present in their atoms or ions. When unpaired electrons can be calculated, with the
a magnetic field is applied, substances which use of spin-only formula. As magnetic moment
are attracted towards the applied magnetic is directly related to number of unpaired
field are called paramagnetic, while the ones electrons, value of m will vary directly with
which are repelled are called diamagnetic. the number of unpaired electrons.

Some substances are attracted very Try this...
strongly and these are called ferromagnetic
substances. What will be the magnetic
moment of transition metal having
Remember... 3 unpaired electrons ?
a. equal to 1.73 BM,
Paramagnetism and b. less than 1.73 BM or
c. more than 1.73 BM ?
ferromagnetism arises due to
Magnetic moments are determined
presence of unpaired electrons in a species. experimentally in solution or in solid state
where the central metal is hydrated or bound
When all electron spins are paired, the to ligands. A slight difference in the calculated
and observed values of magetic moments thus
compound becomes diamagnetic. can be noticed.

Among transition metals Fe, Co, Ni Use your brain power
are ferromagnetic. When magnetic field is A metal ion from the first transition
applied externally all the unpaired electrons in series has two unpaired electrons.
these metals and their compounds align in the Calculate the magnetic moment.
direction of the applied magnetic field. Due to
this the magnetic susceptibility is enhanced. Table 8.7 gives the calculated and
These metals can be magnetized, that is, they observed magnetic moments of cations of 3d
aquire permanent magnetic moment. series.

Try this...

Pick up the paramagnetic species
Cu1⊕, Fe3⊕, Ni2⊕, Zn2⊕, Cd2⊕, Pd2⊕.

171

Table 8.7 Magnetic moments of ions of first transition series elements (values in BM)

Ion Outer electronic Number of Calculated value Experimental

configuration unpaired of magnetic value

electrons moment

Sc3⊕ 3d0 0 0 0

Ti3⊕ 3d1 1 1.73 1.75

V3⊕ 3d2 2 2.84 2.76

Cr3⊕ 3d3 3 3.87 3.86

Cr2⊕ 3d4 4 4.90 4.80

Mn2⊕ 3d5 5 5.92 5.96

Fe2⊕ 3d6 4 4.90 5.3-5.5

Co2⊕ 3d7 3 3.87 4.4-5.2

Ni2⊕ 3d8 2 2.84 2.9-3.0, 4.0

Cu2⊕ 3d9 1 1.73 1.8-2.2

Zn2⊕ 3d10 0 0 0

Problem : Calculate the spin only magnetic Can you tell ?
moment of divalent cation of a transition
metal with atomic number 25. Compounds of s and p block
elements are almost white. What
Solution : For element with atomic number could be the absorbed radiation : uv or
25, electronic configuration for its divalent visible ?
cation will be
The ionic and covalent compounds
3d 4s formed by the transition elements are coloured.
Transition elements contain unpaired electrons
[Ar] in their d orbitals. When the atoms are free or
isolated, the five d orbitals are degenerate;
There are 5 unpaired electrons, so n = 5. or have the same energy. In complexes, the
metal ion is surrounded by solvent molecules
∴ m = 5(5 + 2) = 5.92 BM or ligands. The surrounding molecules affect
the energy of d orbitals and their energies are
Try this... no longer the same [You will learn more about
this in Chapter 9]. As the principal quantum
Calculate the spin only number of ‘d’ orbitals is the same, the amount
magnetic moment of divalent cation of energy required for transition of electron
of element having atomic number 27. from one d orbital to another is quite small.
The small energy required for this transition
In second and third transition series, is available by absorption of radiation
orbital angular moment is significant. having certain wavelength from the visible
Therefore, the simple spin only formula is not region. Remaining light is transmitted and the
useful and more complicated equations have to observed colour of the compound corresponds
be employed to determine magnetic moments. to the complimentary colour of light absorbed.
The magnetic moments further are found to be That means, if red light is absorbed then the
temperature dependent. transmitted light contains excess of other
colours in the spectrum, in particular blue, so
8.6.5 Colour : A substance appears coloured if
it absorbs a portion of visible light. The colour
depends upon the wavelength of absorption in
the visible region of electromagnetic radiation.

172

the compound appears blue. The ions having Thus the colour of a transition metal ion
no unpaired electrons are colorless for example relates to
Cu⊕(3d10); Ti4⊕(3d0). Table 8.8 enlists colours 1. presence of unpaired d electrons
of 3d transition metal ions.
2. d - d transitions
620 nm
3. nature of ligands attached to the metal ion
800 nm
4. geometry of the complex formed by the
400 nm red orange 580 nm metal ion

violet yellow Do you know ?

blue green 560 nm Colour of transition metal ions
may arise due to a charge transfer.
430 nm For example, MnO4 ion has an intense purple
colour in solution. In MnO4 , an electron is
490 nm momentarily transferred from oxygen (O) to
metal, thus momentarily changing O2 to O
Let us see how colour of the transition and reducing the oxidation state of manganese
metal ion depends upon ligand and geometry from +7 to +6. For charge transfer transition
of the complex formed by metal ion. to take place, the energy levels of the two
different atoms involved should be fairly close.
When cobalt chloride (Co2⊕) is dissolved Colours of Cr2O72 , CrO4 , Cu2O and Ni-DMG
in water, it forms a pink solution of the (where DMG = dimethyl glyoxime) complex
complex [Co(H2O)6]2⊕ which has octahedral thus can be explained through charge transfer
geometry. But when this solution is treated transitions.
with concentrated hydrochloric acid, it turns
deep blue. This change is due to the formation 8.6.6 Catalytic Properties : Transition metals
of another complex [CoCl4]2 which has a and their compounds exhibit good catalytic
tetrahedral structure. properties. They have proven to be good
homogeneous and heterogeneous catalysts.
[Co(H2O)6]2⊕ + 4Cl [CoCl4]2 + 6H2O Partly because of their ability to participate in
different oxidation-reduction steps of catalytic
reactions.

Table 8.8 Colour of 3d transition metal ions

Ion Outer electronic configuration Number of unpairedelectrons Colour

Sc3⊕ 3d0 0 Colourless

Ti3⊕ 3d1 1 Purple

Ti4⊕ 3d0 0 Colourless

V3⊕ 3d2 2 Green

Cr3⊕ 3d3 3 violet

Mn2⊕ 3d5 5 Light pink

Mn3⊕ 3d4 4 Violet

Fe2⊕ 3d6 4 Pale green

Fe3⊕ 3d5 5 Yellow

Co2⊕ 3d7 3 Pink

Ni2⊕ 3d8 2 Green

Cu2⊕ 3d9 1 Blue

Cu⊕ 3d10 0 Colourless

Zn2⊕ 3d10 0 Colourless

173

These steps involve changes in the formed are called interstitial compounds.
oxidation states of these metal ions. Sometimes sulphides and oxides are also
Compounds of Fe, Co, Ni, Pd, Pt, Cr, etc. are trapped in the crystal lattice of transition
used as catalysts in a number of reactions. elements. Steel and cast iron are examples of
Their compounds enhance the rate of the interstitial compounds of carbon and iron. Due
chemical reactions. to presence of carbon, the malleability and
ductility of iron is reduced while its tenacity
In homogeneous catalysis reactions, the increases.
metal ions participate by forming unstable
intermediates. In heterogeneous catalysis Some properties of interstitial compounds
reactions on the other hand, the metal provides
a surface for the reactants to react. i. They are hard and good conductors of heat
and electricity.
Examples :
ii. Their chemical properties are similar to the
1. MnO2 acts as a catalyst for decomposition parent metal.
of KClO3.
iii. Their melting points are higher than the
2. In manufacture of ammonia by Haber’s pure metals.
process Mo/Fe is used as a catalyst.
iv. Their densities are less than the parent
3. Co-Th alloy is used in Fischer Tropsch metal.
process in the synthesis of gasoline.
v. The metallic carbides are chemically inert
4. Finely divided Ni, formed by reduction of and extremely hard as diamond.
the heated oxide in hydrogen is an extremely
efficient catalyst in hydrogenation of ethene vi. Hydrides of transition metals are used as
to ethane at 140 0C. powerful reducing agents.

H2C = CH2 + H - H ∆ H3C - CH3 Remember...
Ni, 140 0C • Tungsten carbide is used for

Commercially, hydrogenation with nickel as cutting tools.

catalyst is used to convert inedible oils into • Iron carbide is used in manufacture
of steel.
solid fat for the production of margarine.
8.6.8 Formation of Alloys
5. In the contact process of industrial
Can you recall ?
production of sulfuric acid; sulphur dioxide • What is an alloy ?
• Do atomic radii of 3d transition
and oxygen from the air react reversibly
elements differ largely ?
over a solid catalyst of platinised asbestos.
Transition metals form alloys where
2SO2 + O2 platinised 2SO3 atoms of one metal are distributed randomly
asbestos in the lattice of another metal. The metals with
similar radii and similar properties readily
6. Carbon dioxide and hydrogen are form alloys.

formed by reaction of the carbon monoxide Alloys are classified into ferrous and non-
ferrous.
and steam at about 500 0C with an Fe-Cr

catalyst.

CO + H2O Fe-Cr CO2 + H2
catalyst
(steam)

8.6.7 Formation of interstitial compounds

When small atoms like hydrogen, carbon
or nitrogen are trapped in the interstitial spaces
within the crystal lattice, the compounds

174

Ferrous alloys have atoms of other elements 3MnO2 + 6KOH + KClO3 ∆ 3K2MnO4 +
distributed randomly in atoms of iron in the KCl + 3H2O
mixture. As percentage of iron is more, they
are termed ferrous alloys eg. nickel steel, In neutral or acidic medium the green potassium
chromium steel, stainless steel etc. All steels
have 2% carbon. manganate disproportionates to KMnO4 and
MnO2.
Non-ferrous alloys are formed by mixing
atoms of transition metal other than iron with 3K2MnO4 + 4CO2+ 2H2O 2KMnO4 +
a non transition element. eg. brass, which is MnO2 + 4KHCO3
an alloy of copper and zinc. Ferrous and non-
ferrous alloys are of industrial importance. 3MnO42 + 4H⊕ 2MnO4 +
MnO2 + 2H2O
Uses of alloys
The liquid is filtered through glass
• Bronze, an alloy of copper and tin is tough,
strong and corrosion resistant. It is used for wool or sintered glass and evaporated
making statues, medals and trophies.
until crystallisation occurs. Potassium
• Cupra-nickel, an alloy of copper and nickel
is used for making machinary parts of permanganate forms small crystals which are
marine ships, boats. For example, marine
condenser tubes. almost black in appearance.

• Stainless steels are used in the construction ii. Electrolytic oxidation
of the outer fuselage of ultra-high speed air
craft. In electrolytic oxidation, alkaline solution
of manganate ion is electrolysed between iron
• Nichrome an alloy of nickel and chromium electrodes separated by a diaphragm. Overall
in the ratio 80 : 20 has been developed reaction is as follows :
specifically for gas turbine engines.
2K2MnO4 + H2O + [O] 2KMnO4 + 2KOH
• Titanium alloys withstand stress up to
high temperatures and are used for ultra- The oxygen evolved at the anode converts
high speed flight, fire proof bulkheads and manganate to permanganate.
exhaust shrouds.
The solution is filtered and evaporated to
8.7 Compounds of Mn and Cr (KMnO4 and get deep purple black coloured crystals of
K2Cr2O7) KMnO4.

Remember... 8.7.2 Chemical properties of KMnO4 :
Both KMnO4 and K2Cr2O7 are
strong oxidising agents. a. In acidic medium :

8.7.1 Preparation of potassium permaganate The oxidizing reactions of KMnO4 in acidic
medium
i. Chemical oxidation
i. Oxidation of iodide to iodine :
When a finely divided manganese dioxide
(MnO2) is heated strongly with fused mass 2MnO4 + 10 I + 16H⊕ 2Mn2⊕ +
of caustic potash (KOH) and an oxidising
agent, potassium chlorate (KClO3), dark green 8H2O + I2
potassium manganate, K2MnO4 is formed.
ii. Oxidation of Fe2⊕ to Fe3⊕

MnO4 + 5Fe2⊕ + 8H⊕ 5Fe3⊕ + Mn2⊕
+4H2O

iii. Oxidation of H2S

H2S 2H⊕ + S2

5S2 + 2MnO4 + 16H⊕ 2Mn2⊕ + 5S + 8H2O

175

iv. Oxidation of oxalic acid : Na2Cr2O7 + 2KCl K2Cr2O7 + 2NaCl

2MnO4 + 5H2C2O4 + 6H⊕ 2Mn2⊕ + 8.7.5 Chemical properties of K2Cr2O7 :

10CO2+ 8H2O i. Oxidation of I- from aq. solution of KI

b. In neutral or weakly alkaline medium : by acidified K2Cr2O7 gives I2. Potassium
dichromate is reduced to chromic sulphate.
i. Oxidation of iodide I to iodate IO3 .
Liberated I2 turns the solution brown.
2MnO4 + H2O + I 2MnO2 + 2OH + IO3
K2Cr2O7 + 6KI + 7H2SO4 4K2SO4 +
ii. Thiosulphate (S2O32 ) is oxidised to sulphate Cr2(SO4)3 + 7H2O + 3I2
(SO42 )
ii. When H2S gas is passed through acidified
8MnO4 + 3S2O32 + H2O 8MnO2 + K2Cr2O7 solution, H2S is oxidised to pale
6SO42 + 2OH yellow precipitate of sulphur. Simultaneously

iii. Manganous salt is oxidised to MnO2. potassium dichromate is reduced to chromic

2MnO4 + 3Mn2⊕ + 2H2O 5MnO2 + 4H⊕ sulphate, which is reflected as colour change

8.7.3 Uses of KMnO4 : of solution from orange to green.

• An antiseptic. K2Cr2O7 + 4H2SO4 + 3H2S K2SO4 +
Cr2(SO4)3 + 7H2O + 3S
• For unsaturation test in laboratory.
8.8 Common properties of d block elements

• In volumetric analysis of reducing agents. Physical properties

• For detecting halides in qualitative i. All d block elements are lustrous and
analysis. shining.

• Powerful oxidising agent in laboratory and ii. They are hard and have high density.
industry.
iii. Have high melting and boiling points.
8.7.4 K2Cr2O7 : Preparation of potassium
dichromate iv. Are good electrical and thermal conductors.

In the industrial production, finely powdered v. Have high tensile strength and malleability.
chromite ore (FeOCr2O3) is heated with
anhydrous sodium carbonate (Na2CO3) and a vi. Can form alloys with transition and non
flux of lime in air in a reverbatory furnace. transition elements.

4(FeO.Cr2O3) + 8Na2CO3 + 7O2 ∆ 8Na2CrO4 vii. Many metals and their compounds are
+ 2Fe2O3 + 8CO2. paramagnetic.

viii. Most of the metals are efficient catalysts.

Sodium chromate (Na2CrO4) formed in this Chemical properties
reaction is then extracted with water and
treated with concentrated sulphuric acid to i. All d block elements are electropositive
get sodium dichromate and hydrated sodium metals.
sulphate :
ii. They exhibit variable valencies and form
2Na2CrO4 + H2SO4 Na2Cr2O7 + 2NaCl + colored salts and complexes.
Na2Cr2O7.H2O
iii. They are good reducing agents.
Addition of potassium chloride to concentrated
iv. They form insoluble oxides and hydroxides.
solution of sodium dichromate precipitates
v. Iron, cobalt, copper, molybdenum and
less soluble orange-red coloured potassium zinc are biologically important metals

dichromate, K2Cr2O7. vi. catalyse biological reactions.

176

Differences : Although most properties Table 8.9 : List of minerals and ores of some
exhibited by d block elements are similar, the transition metals.
elements of first row differ from second and
third rows in stabilization of higher oxidation Metals Mineral Ore
states in their compounds. Iron Haematite
Haematite Fe2O3
For example, Mo(V) and W (VI) compounds Copper Magnetite Fe3O4 Chalcopyrite
are more stable than Cr(VI) and Mn (VIII). Zinc Limonite 2Fe2O3, 3H2O Chalcocite
Iron pyrites FeS2 Zinc blende
Highest oxidation state for elements of first row Siderite FeCO3
is +7, and in the case of 3rd row +8 oxidation Chalcopyrite CuFeS2
state as in (RuO4) and (OsO4). Chalcocite
Cuprite Cu2O
Can you recall ?
Zinc blende ZnS
• How are metals found in nature? Zinccite ZnO
Calamine ZnCO3
• Name two the salts of metals that
are found in nature. Metallurgy

Internet my friend Pyrometallurgy Hydrometallurgy Electro-

1. Collect the information on metallurgy
different steps involved in the
extraction of metals from their ores. 8.9.1 Metallurgy : Commercial extraction of

2. Collect information about place where metals from their ores is called metallurgy.
deposits of iron ores are found.
Different methods are used for their extraction
8.9 Extraction of metals
depending on the nature of a metal and its ore.
Most metals are found in the earth’s
crust in the form of their salts, such as a. Pyrometallurgy: A process in which the
carbonates, sulphates, sulphides and oxides. ore is reduced to metal at high temperature
using reducing agents like carbon, hydrogen,
A few metals are nonreactive and occur in aluminium, etc. is called pyrometallurgy.
the free state in the earth’s crust, for example,
silver, gold, and platinum. b. Hydrometallurgy : The process of
extracting metals from the aqueous solution
Mineral : A naturally occuring substance of their salts using suitable reducing agent is
found in the earth’s crust containing inorganic called hydrometallurgy.
salts, solids, siliceous matter etc, is called a
mineral. c. Electrometallurgy : A process in which
metal is extracted by electrolytic reduction of
The mineral which contains high molten (fused) metallic compound is called
percentage of the metal and from which the electrometallurgy.
metal can be extracted economically is called
an ore. Steps Involved in Process of Extraction

Do you know ?

1. Extraction of iron has been
known to Indians since 700 BC.
Indian blacksmiths also knew the
thermo-mechnical process for forging.
During Archaeological studies in Harappa,
Madhya Pradesh different iron objects
belonging to middle iron age were found.
2. Famous iron pillar in Delhi is 1300 years
old and is free of rust till to date.

177

Concentration : After mining the ore from the Iron is extracted from haematite by its
earth’s crust it is subjected to concentration. reduction using coke and limestone. Carbon in
In this step, impurities termed as gangue the limestone is reduced to carbon monoxide.
are removed from the ore and the ore gets Carbon and carbon monoxide together reduce
concentrated. Fe2O3 to metallic iron.

The sand, mud and other unwanted Limestone acts as flux, it combines with
impurities which remain mixed with the ore the gangue material to form molten slag.
deposit are called gangue. Extraction of iron from haematite ore involves
the following steps.
During the process of concentration,
the ore is separated from the gangue material i. Concentration : The powdered ore is washed
using different methods such as washing, in a powerful current of water introduced into
hydraulic classification, magnetic separation, the hydraulic classifier.
froth floatation, etc.
The lighter gangue particles are
The method chosen for concentration separated and the concentrated ore is collected
depends upon the nature of the ore. at the bottom.

8.9.2 Extraction of Iron from Haematite ore ii. Roasting : The concentrated ore is heated
using Blast furnace
in a current of air. The sulfur and arsenic
Do you know ?
Iron is the fourth most impurties present in the ore get converted into
abundant element in the earth’s
crust their oxides and escape as vapour. Ferrous

Composition of Haematite ore : oxide in the ore is converted to Fe2O3.

Fe2O3 + SiO2 + Al2O3 + phosphates 4FeO + O2 2Fe2O3

Gangue The roasted ore is converted into lumps
by sintering.
Internet my friend
Find percentage of oxygen, silicon,{ iii. Reduction (Smelting) : This step is carried
aluminium and iron in earth’s crust. out in a blast furnace. Blast furnace is a tall
cylindrical steel tower which is lined with
refractory bricks.

The height of a typical blast furnace is
25 m and its diameter varies between 5 and
10 m. The furnace works on counter current
principle where the charge comes down and

Iron ore Crushing and grinding Reduction by coke
Reduction by CO
Concentration Reduction Reduction by heat
Reduction by Al
Reduction by electrolysis

Magnetic Gravity Leaching Froth Flotation Refining Pure iron
separation separation

Liquification Distillation Oxidation Electro-refining

178

Burden charging

3Fe2O3 + CO 2Fe3O4 + CO2 2500C Charge distribution
Fe3O4 + CO 3FeO + CO2 Low temperature
CO + N2 distribution
FEO + C Fe + CO Stack
Indirect reduction
MnO + C Mn + CO 12000C
Bosh Swelling and
P2O5 + 5C P2 + 5CO softening
SiO2 + 2C Si + 2CO 17000C
Tuyeres Coke solution
C + O2 Hot blast air 15000C direct reduction
CO2 CO2 + C 2CO 13500C primary slag
Hearth formation
Slag tap hole
1/2 SiO2, 1/3CaO, 1/3Al2O3 Food phenomena
MgO, MnO
Effect of raceway
conditions
Pig iron tap hole Fe, C, Si, S, P,
Mn - 93% Iron

Fig. 8.5 : Blast furnace

hot gases move up the tower. The furnace is Chemical reactions taking place in different
comprised of 3 parts - 1. Hearth, 2. Bosh and zones of the blast furnace
3. Stack
1. Zone of combustion : This is 5 - 10 m from
The charge contaning ore and lime the bottom. The hot air blown through the
stone is introduced into the furnace through a tuyers reacts with coke from the charge to form
cup and cone arrangement. In this arrangement CO.
the cone enables uniform distribution of charge
and the cup prevents the loss of gases. A blast C+ 1 O2 CO, ∆H = -220 kJ
of preheated air is introduced into the furnance 2
below the bosh. The charge and hot air come in
contact with each other and various reactions The reaction is highly exothermic; thus
take place.
the temperature of this zone is around 2000 K.

Some of the CO formed dissociates to form

finely divided carbon.

Reactions in the blast furnace : There are 2 CO 2 C + O2
different temperature zones in the blast furnace.
The temperature goes on increasing from The hot gas rich in CO rises upwards
top to bottom in the furnace. At the top, the
temperature is 500 K. Maximum temperature in the blast furnace. The charge coming down
of the furnace is 2000 K above the tuyers.
There are 3 temperature zones in the furnace. gets heated and reacts with CO. Thus CO acts

as a fuel and also a reducing agent.

2. Zone of Reduction (22-25 m near the top)

1. Zone of combustion - Combustion of coke Here, the temperature is around 900 K.
with O2 in the air.
Fe2O3 is reduced to spongy iron by CO
2. Zone of reduction - Reduction of Fe2O3 to
metallic iron Fe2O3+ 3 CO 2Fe + 3 CO2

3. Zone of slag formation - Formation of slag some amount of Fe2O3 is reduced to iron by
by reaction of gangue with limestone carbon

Fe2O3 + 3C 2 Fe + 3 CO

179

Table 8.9 : Summary of reactions taking place in blast furnace at different temperature zones

Temp K Changes taking place Reaction

500 loss of moisture from ore

900 Reduction of ore by CO Fe2O3 + 3CO 2 Fe + 3CO2

1200 Decomposition of lime CaCO3 ∆ CaO + CO2
1500
Reduction of ore by C Fe2O3 + 3C 2Fe + 3CO
2000 Fusion of iron
Slag formation CaO + SiO2 CaSiO3

12CaO + 2Al2O3 4Ca3AlO3 + 3 O2

3. Zone of slag formation (20 m unit) : 5. Refining : Pure iron can be obtained
The gangue present in the ore is converted by electrolytic refining of impure iron or
to slag. This slag can be used for making other methods given in the flow chart. The
road foundation. Temperature of this zone is choice of extraction technique is governed
1200 K. The gangue contains silica, alumina by the following factors. 1. Nature of ore
and phosphates. A removal of this gangue is 2. Availability and cost of reducing agent,
effected by adding lime-stone in the charge, generally cheap coke is used. 3. Availability of
which acts as flux. Limestone decomposes to hydraulic power. 4. Purity of product (metal)
give CaO (quick lime) required. 5. Value of byproducts for example,
SO2 obtained during roasting of sulphide ores
CaCO3 ∆ CaO + CO2 is vital for manufacture of H2SO4. Knowledge
of electrochemical series provides solutions to
CaO combines with gangue to form many problems.
molten slag of calcium silicate and calcium
aluminate. Commercial forms of Iron

CaO + SiO2 CaSiO3 Iron

12 CaO + 2Al2O3 4Ca3 AlO3 + 3 O2 Cast Wrought Steel

4. Zone of fusion (15 m ht) : MnO2 and Remember...
Ca3(PO4)2 present in the iron ore are reduced
to Mn and P. Some of the silica is also reduced 1. Iron melts at a very high
temperature (1800 K). On
to Si. addition of carbon its melting
point decreases depending upon
The spongy iron coming down in the percentage of carbon.
furnace melt absorbs impurities like C, Si, Mn,
P and S. This molten iron collects at the bottom 2. Mechanical properties of steel can
in the furnace. The slag which is lighter floats be modified by addition of small
on the surface of molten iron. Molten slag and amounts of suitable elements such
iron are collected through separate outlets. as manganese, chromium, sulfur,
nickel etc. These elements are
Molten iron is poured into moulds. called alloying elements and steels
These solid blocks are called pigs. This iron are alloy steels.
contains about 4% of carbon. When pig iron
is remelted, run into moulds and cooled, it
becomes cast iron. The waste gases containing
N2, CO and CO2 escape through the outlet at
the top. These hot gases are used for preheating
the blast of air.

180

Differences between cast iron, wrought iron and (n-1)f orbitals are very similar and are
and steel sensitive to electronic configurations.

Cast iron Wrought Steel Do you know ?

iron Glenn Seaborg first proposed a
revised design of periodic table with
1. Hard and 1. Very soft 1. Neither too a whole new series of elements. When
he showed his design to two prominent
brittle hard nor too inorganic chemists of the time, they warned
him against publishing it. They told him
soft. that tampering with the established periodic
table will affect his career. Seaborg went
2. Contains 2.Contains 2. Contains ahead and published it. He later remarked
“I did not have any scientific reputation, so
4% carbon. less than 0.2 to 2% I published it anyway”. Now we see that
elements 89 -102 (filling of 5f orbitals) fit
0.2% carbon in Seaborg’s proposed order.

carbon 8.11 Properties of f-block elements

3. Used for 3. Used for 3. Used in i. Properties are similar to d block elements

making pipes, making buildings ii. Electrons are added to f subshells of (n-2)
level
manufacturing pipes, infrastructure,
iii. Placed between (n-1)d and ns block
automotive bars for tools, ships, elements

parts, pots, stay bolts, automobiles, Lanthanoids begin with atomic number 57 and
end at 71. Although, historically, lanthanoids
pans, utensils engine bolts weapons etc. are termed as rare earth elements, they are
fairly abundant in earth’s crust. For example,
and rivetts. thulium is found more in abundance than silver
(4.5 x 10-5 vs 0.79 x 10-5 percent by mass). The
Do you know ? name rare earth elements was coined because
Iron possesses a high degree of difficulty in extracting them economically
of magnetism below 1042 K. This is in pure form from other lanthanoids having
known as ferromagnetism. similar chemical properties. Now, due to newer
separation methods like ion exchange resins,
8.10 Inner Transition (f-block) Elements: the separation of these elements has become
Lanthanoids and Actinoids : Elements whose easier and more economical.
f orbitals get filled up by electrons, are called
f block elements. These elements are placed
separately at the bottom of the periodic table.
They are a subset of 6th and 7th periods.

Can you tell ?
• Why f-block elements are

called inner transition metals?
• Are there any similarities between

transition and inner transition metals?

Since f orbital lies much inside the d orbital,
in relation to the transition metals the f block
elements are called inner transition elements.
These elements have 1 to 14 electrons in their
f orbital, 0 or 1 in the penultimate energy level
and 2 electrons in the outermost orbital. The
lanthanoids are characterized by gradual
filling up of 4f and actinides by the 5 f orbitals.
There are 14 elements filling the f orbital in
each series. The relative energies of the nd

181

These metals are soft with moderate decreases. La(OH)3 is the strongest base
densities of about 7 g cm-3. They have high while Lu(OH)3 is the weakest base.
melting (~1000 0C) and boiling points (~3000
0C). Similar to groups 1 and 2, lanthanoids Lanthanoids react with nitrogen and
in the metallic state are very reactive and halogens to give nitrides and halides of the
resemble alkali and alkaline earth metals in formulae LnN and LnX3 repectively. While
their reactivities than transition metals. For doing so, lanthanoids lose their outermost
example, they react with water to give the 3 electrons to form stable compound in +3
metal hydroxide and hydrogen gas oxidation state. When lanthanoids are heated
at elevated temperatures (~ 2800 K) with
2M(s) + 6 H2O (l) 2M(OH)3(s) + 3H2(g) carbon, the carbides with general formula
LnC2 are obtained.
Although, the common oxidation state
for lanthanoids is +3, the +2 oxidation state In +3 oxidation state, many of the
is also important. They all form stable oxides lanthanoids are coloured, mostly green, pink
of the type M2O3 where M is metal ion. Eu2+ and yellow. This is attributed to the electronic
and Yb2+ are the most stable dipositive metal transitions among the f orbitals. Like transition
ions. Higher oxidation states are unusual for metals the electronic spectra of lanthanoids
lanthanoids with the only exception of cerium however, do not get affected with different
which forms a stable +4 species. The energy ligands.
required to break up the metal lattice is heat
of atomization. Lanthanoids have lower heat 8.12 Properties of Lanthanoids
of atomization than transition metals. This
is because with d electrons, transition metals i. Soft metals with silvery white colour and
are much harder and require high heat of moderate densities of ~ 7 g cm-3. Colour and
atomization. Europium and ytterbium have the brightness reduces on exposure to air
lowest enthalpies of vaporization and largest
atomic radii of lanthanoids, resemble barium. ii. Good conductors of heat and electricity.
These two elements resemble alkaline earth
elements; they dissolve in liquid ammonia to iii. Except promethium (Pm), all are non-
give blue conducting solutions. radioactive in nature.

Their ionic radii decrease from 117 pm of iv. The atomic and ionic radii decrease from
La to 100 pm for Lu. This is because 5f orbitals lanthanum (La) to lutetium (Lu). This is known
do not shield the outer 5s and 5p electrons as lanthanoid contraction.
effectively, leading to increase in effective
nuclear charge and decrease in the ionic size. v. Binding to water is common (i.e.) such that
Such large ions have higher coordination H2O is often found in products when isolated
number that varies from 6 (most common) to from aqueous solutions.
9, 10 and upto12 in some cases. For example,
hydrated lanthanum ion is a nonahydrate, vi. Coordination numbers usually are greater
[La(H2O)9]3+. than 6, typically 8, 9,... (up to 12 found).

All the lanthanoids form hydroxides of vii. The lanthanoides are strongly
the general formula Ln(OH)3(Ln represents paramagnetic. Gadolinium becomes
any elements of lanthanoid series). These are ferromagnetic below 16 0C (Curie point). The
ionic and basic. Since the ionic size decreases other heavier lathanoids terbium, dysprosium,
from La3+ to  Lu3+, the basicity of hydroxides holmium, erbium, thulium, and ytterbium
– become ferromagnetic at much lower
temperatures.

viii. Magnetic and optical properties are
largely independent of environment (similar
spectra in gas/solution/solid).

182

Table 8.11 : Electronic configuration and atomic ionic radii of lanthanoids

Electronic configuration Ionic radii
(Ln+3,
Element Symbol Atomic number Expected Observed Atomic
Lanthanum (ground state) (ground radii, pm 6-coordinate),
(+2 oxidation state) (+3 oxidation state) pm
state) 187
183 103
La 57 [Xe]4f 05d16s2 [Xe]4f 05d16s2 4f 0 4f 0 182 102
181 99
Cerium Ce 58 [Xe]4f 26s2 [Xe]4f 15d16s2 4f 1 181 98.3
180 97
Praseodymium Pr 59 [Xe]4f 36s2 [Xe]4f 36s2 4f 2 4f 2 199 95.8
178 94.7
Neodymium Nd 60 [Xe]4f 46s2 [Xe]4f 46s2 177 93.8
176 92.3
Promethium Pm 61 [Xe]4f 56s2 [Xe]4f 56s2 4f 4 175 91.2
174 90.1
183 Samarium Sm 61 [Xe]4f 66s2 [Xe]4f 66s2 4f 5 4f 5 173 89
88
Europium Eu 63 [Xe]4f 76s2 [Xe]4f 76s2 4f 6 86.8
86.1
Gadolinium Gd 64 [Xe]4f 86s2 [Xe]4f 75d16s2 4f 7

Terbium Tb 65 [Xe]4f 96s2 [Xe]4f 96s2 4f 8

Dysprosium Dy 66 [Xe]4f 106s2 [Xe]4f 106s2 4f 9

Holmium Ho 67 [Xe]4f 116s2 [Xe]4f 116s2 4f 10 4f 10

Erbium Er 68 [Xe]4f 126s2 [Xe]4f 126s2

Thulium Tm 69 [Xe]4f 136s2 [Xe]4f 136s2 4f 12 4f 12

Ytterbium Yb 70 [Xe]4f 146s2 [Xe]4f 146s2 4f 13

Lutetium Lu 71 [Xe]4f 145d16s2 [Xe]4f 145d16s2 4f 14

8.12.1 Electronic configuration : The Table 8.12: First, second, third and fourth
electronic configuration of lanthanoids ionization enthalpies of lanthanoids in kJ/mol
is [Xe] 4f0-14 5d 0-2 6s2. This is because
1s22s22p63s23p64s23d104p65s24d105p6 is the Lanthanoid IE1 IE2 IE3
electronic configuration of xenon and we La 538.1 1067 1850.3
can simplify the electronic configuration Ce 528.0 1047 1949
of lanthanoids by [Xe] 4f0-145d0-26s2. The Pr 523.0 1018 2086
lanthanoids involve gradual filling of Nd 530.0 1034 2130
f-orbitals. The energies of 5d and 4f orbitals Pm 536.0 1052 2150
are very close. For lanthanum 4f is slightly Sm 543.0 1068 2260
higher in energy than 5d. The lanthanum has Eu 547.0 1085 2400
electronic configuration [Xe]6s25d1 and not Gd 592.0 1170 1990
[Xe]6s24f1. Gadolinium (Gd) and lutetium Tb 564.0 1112 2110
(Lu) have 5d1 electron to make f-orbital Dy 572.0 1126 2200
half-filled and full-filled which render them Ho 581.0 1139 2200
extra stability. The electronic configuration Er 589.0 1151 2190
of lanthanoids have variable occupancy in 4f Tm 596.7 1163 2284
(0 to 14) orbitals. This can be noticed from Yb 603.4 1175 2415
Table 8.10. Number of electrons in 6s orbitals Lu 523.5 1340 2022
remains constant in the ground state. The
valence shell electronic configuration of these Problem : Which of the following will
elements, thus can be represented as: (n-2)f 0,2- have highest forth ionization enthalpy IE4?
14 (n-1)d0,1,2 ns2 La4⊕, Gd4⊕, Lu4⊕.

The electronic distribution in different Solution : First write electronic
orbitals of elements in their ground and excited configuration of that element/ion. Check
states are shown in Table 8.11. for any unpaired electrons present. The
energy required for removal of that electron
Try this... will be less as compared to the energy
required to remove an electron from an
Fill in the blanks in Table 8.11. electron pair.Also compare the energies of
the orbitals occupying those electrons. It
Ionization Enthalpies will be easier to remove an electron from
an orbital that is lower in energy than the
Can you recall ? one with higher in energy. First ionization
enthalpy generally decrease across the
• What is ionization enthalpy? period.
• Some elements have variable

oxidation states and some have only
two. Can this be justified based on
their ionization enthalpies?

The ionization enthalpies of lanthanoids are
given in Table 8.12

184

8.12.2 Oxidation state : +3 oxidation state is 4f subshell. Thus the 4f electrons shield each
common to all elements in which 2 electrons other from the nuclear charge poorly owing to
of s- subshell and one from d or f- subshell their diffused nature. With increasing atomic
are removed. The 4f electrons are strongly number and nuclear charge, the effective
screened by inner electrons of 5s and 5p nuclear charge experienced by each 4f
electrons. Thus, 4f electrons are not involved in electrons increases. As a result, the whole of
the bonding. Besides these, some lanthanoids 4f electron shell contracts at each successive
show oxidation states +2 and +4. They are element.
formed in case of f0, f7 , f14 configurations or
resulting ions. Fig. 8.6 : Ionic radii of lanthanoids in +3
oxidation state
For example : Ce4⊕ (f 0) ; Eu2⊕ and Tb4⊕ (f 7) ;
In section 8.3.4 we have learnt about
Yb2⊕ (f 14) Refer Table 8.11. the magnetic behaviour of transition metal
complexes.
8.12.3 Colour and Spectra :
Use your brain power
Some trivalent ions (M3⊕) are coloured in
solid state as well as in solution. The colour of • Do you think that lanthanoid
lanthanoid ion is due to f-f transitions which
correspond to energy in the visible region of complexes would show
the electromagnetic spectrum. The colour of magnetism?
ions having nf electrons is about the same as
those having (14 – n)f electrons. (where n is an • Can you calculate the spin only
integer 1-13).
magnetic moment of lanthanoid
Ln No. of Colour (14 -n) complexes using the same formula
ion f-electrons green f-electrons that you used for transition metal
Pr3⊕ =14-2 =12 complexes?
4f 2 green
Tm3⊕ pink n f-electrons • Calculate the spin only magnetic
Nd3⊕ 4f 12 =12
4f 3 pink moment of La3+. Compare the value
Er3⊕ (14 -n) with that given in Table 8.13 Is it same
4f 11 f-electrons or different?
=14-3 =11

nf-electrons
=11

8.12.4 Atomic and ionic radii (Lanthanoid
Contraction) : As we move along the
lanthanoid series, there is a decrease in
atomic and ionic radii (Fig.8.6). This steady
decrease in the atomic and ionic radii is called
Lanthanoide contraction. As we move from
one element to another the nuclear  charge
increases by one unit and one electron is added.
The new electrons are added to the same inner

185

Table 8.13: Effective magnetic moments of For example, mixed oxide of europium and
lanthanoids in +3 oxidation state yttrium (Eu,Y)2O3 releases an intense red
colour when bombarded with the high energy
Ln Ln3+ No. of Observed electrons.The optoelectronics applications use
lanthanoid ions as active ions in luminescent
oxidation unpaired magnetic materials. The most notable application is the
Nd: YAG laser (Nd: YAG = neodymium doped
state electrons moment, yttrium aluminium garnet). Erbium-doped fibre
amplifiers are significant devices in the optical-
μeff B.M fibre communication systems. Lanthanoids
are used in hybrid cars, superconductors and
La 4f0 0 0 permanent magnets.

Ce 4f1 1 2.3-2.5 8.14 Actinoids : The last row of elements in the
periodic table is the actinoid series. It begins
Pr 4f2 2 3.4-3.6 at thorium (Z =72) and ends at lawrencium
(Z=103). Most of these elements are not
Nd 4f3 3 3.5-3.6 found in nature. They are all radioactive and
man-made. The half-lives of the isotopes
Pm 4f4 4 -- of thorium (Th-232=1.4 x 1010 years) and
uranium (U-238=4.5 x 109 years) are so long
Sm 4f5 5 1.4-1.7 that these elements exist in rocks on earth.
The long lived isotopes such as thorium,
Eu 4f6 6 3.3-3.5 protactinium, uranium, neptunium, plutonium
and americium, are studied in more details.
Gd 4f7 7 7.9-8.0 These elements have high densities (~ 15-20
g cm-3), high melting points (~1000 oC) and
Tb 4f8 6 9.5-9.8 high boiling points (~3000 oC). Actinoids are
less reactive than lanthanoids. For example,
Dy 4f9 5 10.4-10.6 they react with hot, but not cold water to
give the hydroxide and hydrogen gas. Unlike
Ho 4f10 4 10.4-10.7 lanthanoids, they exhibit a range of oxidation
numbers in their compounds which varies
Er 4f11 3 9.4-9.6 from +2 to +8. The most common oxidation
numbers of the actinoids are shown in Fig. 8.7.
Tm 4f12 2 7.1-7.6
As can be seen from Fig.8.7, the most
Yb 4f13 1 4.3-4.9 common oxidation state of early actinoids
reflects the loss of all outer electrons which
Lu 4f14 0 0 is similar to transition metals than the
lanthanoids. A ready loss of 5f electrons by
8.13 Applications early actinoids indicates that these electrons
are much closer in energy to 7s and 6d electrons
The lanthanoid compounds are present than the 4f electrons to 6s and 5d electrons as
in every household. It is inside the colour in lanthanoids. All three sets of orbitals that
television tubes. When electrons are bombarded is 6d, 5f and 7s have similar energies. For Th,
on certain mixed lanthanoid compounds, they Pa and Np difference in energy levels is small
emit visible light over a small wavelength
range. Therefore, the inside surface of a
television tube or computer monitor is coated
with tiny patches of three different lanthanoid
compositions to give three colours that make
the colour image.

Promethium Europium Actinium Brekelium
Nobelium Mendelenium Sebargium

186

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With the help of internet find out the applications of the elements listed in
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Element Applications

Lanthanum

Ytterbium

Erbium

Praseodymium

Samarium

Promethium

so electrons occupy either 6d or 5f oribitals. In attraction by 7s electrons. As a result all three
actinoids series the 5f are orbitals appreciablly sets of orbitals i.e. 6d, 5f and 7s have very
lower in energy, thus from Pu onwards 5f shell similar energies. The ionic radius decreases
gets filled in a regular way. as we move across the series which is known
as ‘Actinoid contraction’. This is attributed to
The electronic configuration of actinoids is poor shielding offered by f electrons.
[Rn] 5f0-14 6d0-2 7s2, where Rn is the electronic
configuration of radon. As seen from Fig. 8.7, Fig. 8. 8 Figure depicting contraction of ionic
the most stable oxidation state in actinoids radii of lanthanoides and actinoids
is +3. The highest common oxidation states
of early actinoids reflect the loss of all outer 8.15 Properties of Actinoids
electrons which is similar to transition metals i. Similar to lanthanoids, they appear silvery
than lanthanoids. For example, uranium has white in colour.
electronic configuration of [Rn]7s25f 36d1. The ii. These are highly reactive radioactive
formation of +6 oxidation state corresponds elements
to an electronic configuration of [Rn]. Similar
to lanthanoids, loss of s and d electrons occur
before f electrons, in formation of 3+ ions. A
ready loss of 5f electrons by early actinoids
indicates that these electrons are much closer
in energy to 7s and 6d electrons than the 4f
electrons are to 6s and 5d electrons in the
lanthanoids. In turn, 5f and 6s orbitals expand
as they are partially shielded from the nuclear

Oxidation numbers 6
5
4
3
2
1

Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

Fig. 8. 7 The most common oxidation numbers of actinoids

187

iii. Except promethium (Pm), all are non- Do you know ?
radioactive in nature
Uranium is another actinoid which
iv. They experience decrease in the atomic and
ionic radii from actinium (Ac) to lawrencium is in great demand as it is used in
(Lw), known as actinoid contraction
the nuclear reactors. One of the extraction
v. They usually exhibit +3 oxidation state.
Elements of first half of the series usually methods for uranium has a very interesting
exhibit higher oxidation states.
chemistry. The ore containing U(IV) oxide,
8.16 Applications of actinoids : We have
seen that the half-lives of natural thorium UO2, is first treated with Fe(III) ion to give
and uranium isotopes are so long that we get U(VI) oxide, UO3
very negligible radiation from these elements.
We find them in everyday use. For example, UO2 (s) + H2O(l)  UO3 (s) + 2H⊕ (aq)
Th(IV) oxide, ThO2 with 1% CeO2 was used + 2e
as a major source of indoor lighting before
incandescent lamps came into existence only Fe3⊕ (aq) + e  Fe2⊕ (aq)
because these oxides convert heat energy from
burning natural gas to an intense light. Even Addition of H2SO4 to this solution produces
today, there is a great demand for these lights uranyl sulphate containing UO22⊕ cation:
for outdoor camping.
UO3 (s) + H2SO4 (aq)  UO2SO4 (aq) +
H2O (l)

After purification, ammonia is added to the

solution giving bright yellow precipitate of

ammonium diuranate, (NH4)2U7O7:
2 UO2SO4 (aq) + 6 NH3 (aq) + 3 H2O (l)
 (NH4)2U7O7 (s) + 2 (NH4)2SO4 (aq)
This yellow cake is the marketable form of

uranium!

Similarities and differences between lanthanides and actinoids

Similarities Differences

Both the series show a +3 oxidation state Lanthanoids show a maximum oxidation state of +4
while actinoids show oxidation states of +3, +4, +5,
+6 and +7

In both the series, the f-orbitals are filled Lanthanoids do not form complexes easily. Actinoids
gradually have a greater tendency to form complexes with
ligands such as thio-ethers

Ionic radius of the elements in both series All lanthanoids are non-radioactive except
decreases with an increase in atomic number promethium but actinides are radioactive in nature

The electronegativity of all the elements in Lanthanoids do not form oxocations, but actinides
both the series is low and are said to be highly form oxocations such as UO+, PuO+, NpO2+
reactive

The nitrates, perchlorates and sulphates of all Most of the lanthanoids are colourless in nature
the elements are soluble while the hydroxides, whereas the actinoids are coloured ions
fluorides and carbonates are insoluble

188

Table 8.14 Electronic configuration of actinoids and their ionic radii in +3 oxidation state

Element Symbol Atomic Electronic configuration *Atomic *Ionic radii
number radii, pm (Ac3⊕), pm

ground state +3
oxidation

state

Actinium Ac 89 [Rn]5f 06d17s2 5f 0 203 126

Thorium Th 90 [Rn]5f 06d27s2 5f 1 180 -

Protactinium Pa 91 [Rn]5f 26d17s2 5f 2 162 118

Uranium U 92 [Rn]5f 36d17s2 5f 3 153 118

Neptunium Np 93 [Rn]5f 46d17s2 5f 4 150 116

Plutonium Pu 94 [Rn]5f 66d07s2 5f 5 162 115

Americium Am 95 [Rn]5f 76d07s2 5f 6 173 114

Curium Cm 96 [Rn]5f 76d17s2 5f 7 174 112

Berkelium Bk 97 [Rn]5f 96d07s2 5f 8 170 110

Californium Cf 98 [Rn]5f 106d07s2 5f 9 186 109

Einsteinium Es 99 [Rn]5f 116d07s2 5f 10 186 98

Fermium Fm 100 [Rn]5f 126d07s2 5f 11 198 91

Mendelevium Md 101 [Rn]5f 136d07s2 5f 12 194 90

Nobelium No 102 [Rn]5f146d07s2 5f 13 197 95

Lawrencium Lr 103 [Rn]5f 146d17s2 5f 14 171 88

Table 8.15 : Some comparison between Pre-Transition, Lanthanoids and Transition Metals

Pre-Transition Metals Lanthanoids Transition Metals

Essentially monovalent - Essentially in (+3) Show variable oxidation
show group (n+) oxidation oxidation state states
state (+2/+4 for certain
configurations)

Periodic trends dominated Lanthanoid contraction of Size changes of Mn⊕, less
by effective nuclear charge Ln3⊕ marked
at noble gas configuration

Similar properties for a Similar properties Substantial changes in
given group properties

Always 'hard' (O, X, Always 'hard' (O, X, heavier metals
N donors, preferably N donors, preferably (increasingly from Fe-Cu)
negatively charged) negatively charged) may show a 'soft' character

No ligand field effects Insignificant ligand field Substantial ligand field
effects effects

Poor coordination High coordination Coordination number 6
properties(C.N. determined numbers (C.N. determined is typical maximum (many
by size) by size) exceptions)

Flexibility in geometry Flexibility in geometry Fixed geometries (ligand
field effects)

No magnetism Show magnetism Show magnetism

189

8.17 Postactinoid elements : You have seen With half-lives of miliseconds only a little is
that elements with atomic number greater than known about the chemistry of these elements.
92 are called ‘Transuranium’. Elements from Rutherfordium forms a chloride, RfCl4, similar
atomic number 93 to 103 now are included to zirconium and hafnium in the +4 oxidation
in actinoid series and those from 104 to 118 state. Dubnium resembles to both, group 5
are called as postactinoid elements. The transition metal, niobium(V) and actinoid,
postactinoid elements that are known so far protactinium(V).
are transition metals. They are included as
postactinoids because similar to actinoid Do you know ?
elements, they can be synthesized in the Traditionally, no element was
nuclear reactions. So far, nine postactinoid named after a still-living scientist.
elements have been synthesized. It is difficult This principle was put to an end with
to study their chemistry owing their short half- naming the element 106 as ‘Seaborgium’.
lives. For example, element 112 has a half-life
of only 2.8 x 10-4 seconds.

Exercises

1. Choose the most correct option.

i. Which one of the following is c. 3d10, 4s1; 3d10, 4s0
dimagnetic
d. 3d8, 4s1; 3d10, 4s0

a. Cr2⊕ b. Fe3⊕ vii. Which of the following have d0s0
configuration
c. Cu2⊕ d. Sc3⊕

ii. Most stable oxidation state of a. Sc3⊕ b. Ti4⊕
Titanium is
c. V5⊕ d. all of the above

a. +2 b. +3 viii. Magnetic moment of a metal
complex is 5.9 B.M. Number of
c. +4 d. +5 unpaired electrons in the complex is

iii. Components of Nichrome alloy are a. 2 b. 3
are

a. Ni, Cr, Fe b. Ni, Cr, Fe, C c. 4 d. 5

c. Ni, Cr d. Cu, Fe ix. In which of the following series
all the elements are radioactive in
iv. Most stable oxidation state of nature
Ruthenium is

a. +2 b. +4 a. Lanthanides

c. +8 d. +6 b. Actinides

v. Stable oxidation states for chromiom c. d-block elements
are
d. s-block elements

a. +2, +3 b. +3, +4 x. Which of the following sets of ions
contain only paramagnetic ions
c. +4, +5 d. +3, +6

vi. Electronic configuration of Cu and a. Sm3⊕, Ho3⊕, Lu3⊕
Cu+1 b. La3⊕, Ce3⊕, Sm3⊕

a. 3d10, 4s0; 3d9, 4s0 c. La3⊕, Eu3⊕, Gd3⊕
d. Ce3⊕, Eu3⊕, Yb3⊕
b. 3d9, 4s1; 3d94s0

190

xi. Which actinoid, other than uranium, ii. Name different zones in the Blast
occur in significant amount furnace. Write the reactions taking
naturally? place in them.

a. Thorium b. Actinium iii. What are the differences between cast
iron, wrought iron and steel.
c. Protactinium d. Plutonium
iv. Iron exhibits +2 and +3 oxidation states.
xii. The flux added during extraction of Write their electronic configuration.
Iron from teamatite are its? Which will be more stable ? Why ?

a. Silica v. Give the similarites and differences in
elements of 3d, 4d and 5d series.
b. Calcium carbonate
vi. Explain trends in ionisation enthalpies
c. Sodium carbonate of d block elements.

d. Alumina vii. What is meant by diamagnetic and
paramagnetic? Give one example
2. Answer the following of diamagnetic and paramagnetic
transition metal and lanthanoid metal.
i What is the oxidation state of
Manganese in (i) MnO42- (ii) MnO4- ? viii.Why the ground-state electronic
configurations of gadolinium and
ii. Give uses of KMnO4 lawrentium are different than expected?
iii. Why salts of Sc3⊕, Ti4⊕, V5⊕ are
ix. Write steps involved in metallugical
colourless ? process

iv. Which steps are involved in x. Cerium and Terbium behaves as good
manufacture of potassium dichromate oxidising agents in +4 oxidation state.
from chromite ore ? Explain.

v. Balance the following equation xi. Europium and xtterbium behave as
good reducing agents in +2 oxidation
(i) KMnO4 + H2C2O4 MnSO4 + state explain.
K2SO4 + H2O + O2
Activity :
(ii) K2Cr2O7 + KI + H2SO4 K2SO4 +
Cr2(SO4)3 + 7H2O + 3I2 Make groups and each group
prepares a powerpoint presentation
vi. What are the stable oxidation states on properties and applications of one
element. You can use your imagination
of plutonium, cerium, manganese, to create some innovative ways of
presenting data.
Europium ?
You can use pictures, images, flow
vii. Write probable electronic configuration charts, etc. to make the presentation
of chromium and copper. easier to understand. Don’t forget to
cite the reference(s) from where data
viii. Why nobelium is the only actinoid for presentation is collected (including
with +2 oxidation state? figures and charts). Have fun!

ix. Explain with the help of balanced
chemical equation, why the solution of
Ce(IV) is acidic.

x. What is meant by ‘shielding of
electrons’ in an atom?

xi. The atomic number of an element is
90. Is this element diamagnetic or
paramagnetic?

3. Answer the following

i. Explain the trends in atomic radii of d
block elements

191