www.pdfgrip.com

BASIC CONCEPTS OF
INORGANIC CHEMISTRY
(Second Edition)

D. N. Singh
(M.Sc., Ph.D)
Formerly, Reader in Chemistry
P.G.M.S. College, Motihari.
(B.R. Ambedkar Bihar University Muzaffarpur)

&KDQGLJDUK‡'HOKL‡&KHQQDL


www.pdfgrip.com

The aim of this publication is to supply information taken from sources believed to be valid and reliable. This is not an
attempt to render any type of professional advice or analysis, nor is it to be treated as such. While much care has been
taken to ensure the veracity and currency of the information presented within, neither the publisher nor its authors bear
any responsibility for any damage arising from inadvertent omissions, negligence or inaccuracies (typographical or
factual) that may have found their way into this book.
Copyright © 2012 Dorling Kindersley (India) Pvt. Ltd
Licensees of Pearson Education in South Asia
No part of this eBook may be used or reproduced in any manner whatsoever without the publisher’s prior written
consent.
This eBook may or may not include all assets that were part of the print version. The publisher reserves the right to
remove any material present in this eBook at any time.
ISBN 9788131768617
eISBN 9788131798683

Head Office: A-8(A), Sector 62, Knowledge Boulevard, 7th Floor, NOIDA 201 309, India
Registered Office: 11 Local Shopping Centre, Panchsheel Park, New Delhi 110 017, India


www.pdfgrip.com

DEDICATED TO
“MY GURUS”


www.pdfgrip.com

This page is intentionally left blank.


www.pdfgrip.com

CONTENTS
Preface
  1.Periodic Table and 
Periodicity of Properties

xi
1

Mendeleev’s periodic law
1
Modern Periodic Table
2
IUPAC table

2
Periodic table and Aufbau principle
3
Determination of group and 
3
period of an element
Classification of elements
5
Atomicity and elements of the periodic table 8
Effective nuclear charge
9
Atomic and ionic size
10
Ionization energy
17
Electron affinity
19
Electronegativity
21
Physical properties and the periodic table 23
Artificial elements
29
Some records of periodic table
30
  2.Chemical Bonding and
Molecular Structure 
Different types of chemical bonds
Summery of bond types
Ionic bond
Born – Haber cycle

Cations of stable electron configuration
Polarization and its effects
Lattice energy
General properties of ionic compounds
Covalent bonding
σ and π bonds
Comparison of σ and π bond
Electronegativity
Coordinate covalent bond
Electronegativity and Dipole moment
Lewis structure (or Dot structure)
Resonance structure
Hybridization
Resonance
Resonance effect
Resonance energy

33
33
34
34
37
38
38
40
42
42
43
44
46

47
48
52
54
55
61
63
63

Bond length
Bond energy
VSEPR model
Structure and shape
Shape of Molecules
Bond angle
Odd electron molecules
Molecular orbital model
Shape and symmetry of molecular orbitals
Formation of π bonds
Diatomics of the first period elements
Homonuclear Diatomics of Second Period
Elements
NO molecule
CO molecule
Hydrogen bond
Van der Waals, forces
Metallic bond
Metal structures
Electron gas model
Band model

Cohesive energy
  3. Acids and Bases
Bronsted–Lowry theory
Strength of acids and bases
PH
Buffer solution
Henderson equation
Amphoterism
Acid strength
Strength of hydra acids
Strength of Oxy acids
Base strength
Strength of hydra bases
Strength of hydroxide bases
Lewis Acid – Base theory
Hard and soft acids and bases
  4. Chemical Reaction
Types of chemical reactions
Hydrolysis

64
66
68
68
68
72
76
76
78
79

80
81
87
88
89
94
95
95
96
97
98
102
102
103
105
106
106
107
109
109
110
113
114
115
116
118
122
123
126



www.pdfgrip.com

Oxidation–reduction reactions
Oxidation number
Oxidizing agents
Reducing agents
Oxidizing and reducing agents
Strength of oxidants and reductants
Equivalent weights of oxidizing
and reducing agents
Balancing of redox reactions
Oxidation number method
Ion–electron method
Some important half reactions
  5. Transition Elements 
General properties
Atomic and ionic radii
Bonding in Transition metals and
its effect on properties
Electrode potential
Oxidation states
Paramagnetic nature of transition metal
compounds
Complex compound formation
Colour of transition metal compounds
The d – d transition
Colour and charge transfer
Hydrolysis of transition metal compounds
Catalytic property

Lanthanides and actinides
  6. Coordination Chemistry
Coordination number
Types of ligands
Chelates
Conditions for complex formation
Werner’s coordination theory
Nomenclature
Valence bond model for complexes
Crystal field model of bonding
Stability and CF model
Magnetic properties and CF model
Colour and CF model
Isomerism in coordination compounds
Geometrical isomerism
Optical isomerism
Preparations of complex compounds
Stability of complex 
compound in a solution
Applicability of complex compounds

129
130
132
132
133
134
136
137
137

137
141
144
146
146
147
148
148
151
152
152
152
154
154
155
160
166
166
166
167
167
168
171
172
176
178
180
181
183
185

188
191
191
192

Organometallic compounds
Preparation of organometallies
Bonding in organometallic compounds
Bonding in alkene complexes
  7. Abundance and metallurgy
Chemical elements in the Earth’s crust
Cosmic abundance of elements
Abundance in oceans
Occurrence of metals
Metallurgy
Terms used in metallurgical process
Concentration of ore
Leaching
Isolation of metal from concentrated ore
Thermal (or chemical reduction)
Auto reduction
Electrolytic method of reduction
Displacement of one metal by the other
Purification of isolated metals
Sodium
Magnesium
Calcium
Aluminium
Iron
Ashoka Pillar at Delhi

Pig iron
Grey cast iron
White cast iron
Wrought iron
Steel
Special steel
Conversion of iron into steel
Steel from wrought iron
Comparison Pig iron, Wrought and Steel
Heat treatments of steel
Chemically pure iron
Brief chemistry of iron
Corrosion of iron
Important compounds of iron
Copper
Alloys of copper
Brief chemistry of Cu
Important compounds and complexes
Zinc
Brief chemistry of Zn
Important compounds
Mercury
Brief chemistry of Hg

192
194
195
196
201
201

201
202
202
203
203
205
207
207
208
208
209
209
209
211
212
214
215
218
218
220
220
220
220
221
221
221
223
224
224
225

225
228
228
230
232
233
235
236
237
239
239
240


www.pdfgrip.com

Important compounds
Tin
Allotropic forms
Chemical reaction 
Lead
Physical properties
Important compounds
  8. Important Chemical Compounds
Metal compounds
NaOH
Na2Co3

K2Cr2O7
KMnO4

Non-metal compounds
O3
H2O2
NA2S2O3
H2S
  9. Hydrogen and Its Chemistry

242
243
244
245
246
247
248
252
252
255
258
261
265
266
266
268
273
274
278

Position in the periodic table
Isotopes of H
Oxidation states and bonding

Preparation of H2
Uyeno’s reaction
Bosch’s process
Lane’s process
Laboratory Preparation of H2
Nascent Hydrogen
Compounds of protium (H)
Water
Zeolite water
Water clathrates
Hard and soft water
Temporary hardness
Permanent hardness
Inorganic exchangers
Organic exchangers
Bad effects of hard water
Structure of water and ice
Density of water and ice
Density of water at 4oC
Heavy water

278
278
280
281
282
283
283
284
284

285
286
289
289
289
290
290
291
291
291
292
293
293
293

10. Group – 1(IA) The alkali metals

299

Chemical reactions
OXO salts
Halides
Flame colour

301
302
303
305

Alkali metals and liquid NH3

Anomalous behaviour of Li

305
306

11. Group – 2(IIA) [Be, Mg, Ca, Sr, Ba, Ra] 309
Properties which decrease down the group
Properties which increase down the gorup
Oxidation states and nature of bond
Hydrides
Halides
Oxides and hydroxides
Oxo salts
Flame colouration
12. Group – 11(IIB) Cu, Ag, Au
Metallic bond strength
Sublimation energy
Atomic and ionic radii
Ionization energy
Noble metal nature
Malleability, thermal and electrical
conductivities
Oxidation states
Magnetic properties
Colour of compounds
Solubility of Silver–Hlides
Chemistry of photography
13. Group – 12(IIB) Zn, Cd, Hg

310

311
311
311
312
314
316
318
321
322
322
322
322
322
323
323
324
325
326
326
329

Ionization energy
Oxidation states
Nature of bonds
Electrode potential
Magnetic properties
Colour of compounds
Some useful compounds
Biochemistry of Zn, Cd and Hg


330
330
331
331
332
332
333
333

14. Group – 13 (IIIA) B, Al, Ga, In, Td

336

Oxidation states and nature of bond
Hydrides
Diborane
Structure of B2H6
Borazole
Boric acid
Halides
Lewis acid strength of BX3
Alums
Isolation of B
Crystalline B

337
339
339
340
342

345
346
347
347
348
348


www.pdfgrip.com

15. Group – 14(IVA) C, Si, Ge, Sn, Pb
Catenation
Allotropy and structure
Graphite
Diamond
Fullerenes
Semiconductor property of Si and Ge
Physical properties of group – 14 elements
Oxidation states and bonding
Carbides
Oxides
Cyanogens
HCN
Cyanides
Halides
Hydrides
Silicones
Silicates
Isolation of Si
16. Group – 15(VA)N, P, As, Sb, Bi

Allotropes of P
Oxidation state and nature of bond
Hydrides
NH3
PH3
Oxides on N and P
N2O
NO
N2O3
NO2
N2O5
P4O6 and P4O10
HNO2
HNO3
Aquaregia
H3PO2
H3PO3
Phosphoric acids
Acid strength of H3PO2, H3PO3 and P3PO4
Halides
Isolation of N and P
Fertilizers

351
351
352
352
353
353
354

355
355
356
357
360
360
361
362
364
364
366
368
372
373
374
375
377
378
378
379
379
380
380
381
381
382
383
385
386
386

387
389
389
392
393

17. Group–16(VIA) O, S, Se, Te, Po

399

Physical state of the elements
Allotropy of O and S
Effect of heat on S

399
400
401

Viscosity of liquid S and temperature
Oxidation state and nature of bond
Hydrides
H2O2
Strength of H2O2
Acid strength of H2O2 and H2O
Structure of H2O2
Halides
SOCl2
Oxides
Oxo acids
18.Group – 17 (VIIIA) Halogens

F, Cl, Br, I, and At
Physical state
Special properties of F
Oxidation state and bonding
Formation of X2
Manufacture of Cl2
Manufacture of Br2
Manufacture of I2
Reactions of X2
Hydrogen halides
HF
HCl
HBr 
Hl
Halides
Preparation of anhydrous halides
Halogen oxides
Oxo acids
Acid Strength
Oxidizing power
CIO–n anions
Halic acids
Perhalic acids
Interhalogen compounds
Pseudohalogens and pseudohalides

401
401
403
404

405
405
406
406
406
408
414
420
420
422
422
423
424
425
425
426
427
428
429
430
430
431
433
433
435
436
436
437
439
440

441
443

19. Group – 18 The Noble Gases

447

Atomicity
Radii
Water solubility
Special properties of He
Uses of noble gas
Clathrate
Xe compounds
Structure of Xe – Compounds

447
448
448
448
449
449
449
452


www.pdfgrip.com

20. Analytical Chemistry
Carbonates

Sulphite
Nitrite
Chlorides
Bromides
Iodide
Nitrate

454
455
456
459
460
461
462
463

Sulphate
Tests for basic radicals
Flame test
Borax bead test
Solution test for basic radicals
Test of NH+4 ion

464
464
466
466
468
469


21. Problems on Inorganic Reactions

484

Additional Practice Questions

491


www.pdfgrip.com

This page is intentionally left blank.


www.pdfgrip.com

Preface

Inorganic Chemistry is the least taught portion of Class 12 Chemistry course.
Good part about the questions from the inorganic portion is that they are solely based on the principles. But
the books already available in the market talk least of the principles and more about the experimental details
and other things. This book mainly deals with the principles of inorganic chemistry and uses experimental
conditions only when it is needed.
This book has been written to meet the needs of students from different courses and different boards like
CBSE, ICSE and various State boards.
I am indebted to several of my colleagues for their suggestion and criticism. I am also thankful to my family
members for encouragement, especially my daughter, Shiva, in the final draft.
I sincerely appreciate the help of Mr Lalit Kumar Gupta in the preparation of the manuscript.
I hope the readers will appreciate this book and any comments/suggestions towards improving the text would
be welcome.


D. N. Singh


www.pdfgrip.com

This page is intentionally left blank.


www.pdfgrip.com

1

Periodic Table and
Periodicity of Properties

In 19th century, (before 1869) many pseudoscientific arrangement of elements were proposed. But none
withstood the test of time and so were rejected. During 1869, first of all, Mendeleev (Russian chemist)
realized that atomic weights of elements are related with their properties. On this observation, he proposed a
law for the arrangement of elements.

Mendeleev’s Periodic Law
This law states that properties of elements and their compounds are the periodic function of their atomic
weights. He, then arranged the elements in increasing order of their atomic weights. It resulted into a tabular
form, so it is called Mendeleev’s periodic table. Only 63 elements were known when this classification was
presented.

Important features of the Table
(i) Mendeleev’s Periodic table forms vertical columns and horizontal rows for elements.
(ii) Vertical columns are called groups. There are eight groups in the table (Group I to Group VIII). Elements

of a group are similar in properties.
(iii) All groups from I to VII are divided into sub-groups A and B (i.e., IA, IB, IIA, IIB etc.). Group VIII has no
sub-groups.
(iv) Horizontal rows are known as periods. There are seven periods (1 to 7).
(v) Elements of the same period differ in properties.

2 – Periods


Li Be
Metals

B  C  N  O  F
Non-metals

(vi) Sub-group elements also differ in properties.
IA Na (highly reactive metal)
IB Cu (noble metal)

Merits of Mendeleev’s Periodic Table
(i) Correction of atomic weights: Many doubtful atomic weights were corrected using Mendeleev’s periodic
table, Example, Be.
(ii) Vacant places for undiscovered elements: Mendeleev left vacant places in his table for elements to be
discovered, (Ga, Ge etc.). Not only that, he also predicted properties of those elements which were
found true when the elements were actually discovered. For example, Ga and Ge were not known


www.pdfgrip.com
2  Basic Concepts of Inorganic Chemistry
when Mendeleev proposed his periodic table. He named these elements EKAAluminium (Ga) and

EKASILICON (Ge) because he believed that they would be similar to Al and Si. Mendeleev’s predictions
were found to be true.

Defects of Mendeleev’s Periodic Table
(i) H has got two places in Mendeleev’s periodic table, in Group I (alkali metals) and Group VII
(halogens).
(ii) Some elements are in reverse order of atomic weights like Ar (39.9)–K (39.1), Co (58.9)–Ni (58.7), Te
(127.6)–I (126.9).
(iii)Highly reactive alkali metals (Li to Cs) are with noble coinage metals (Cu, Ag, Au).
(iv) Gr VIII has three elements at a place (Fe, CO, Ni), (Ru, Rh, Pd) and (Os, Ir, Pt), instead of one
element.
(v) Metals like Mn, Tc, Re are placed with non-metals and halogens (F to l ) in Group VII.
(vi) Isotopes have got no places in Mendeleev’s periodic table.
(vii) There are 15 lanthanides in the same Group III.

Modern Periodic Table
Mosley (1914), on the basis of his experiment observed that atomic number is the fundamental property of
an element. This concept changed the basis of Periodic Table from atomic weight to atomic number. Mosley
proposed his own periodic law, known as Modern Periodic Law. It states “Properties of elements and their
compounds are the periodic function of their atomic numbers”. Elements were, then, arranged in increasing
order of their atomic numbers. The resultant table is known as Modern Periodic Table. In this table sub
groups are separated and so the table becomes longer in size. Therefore, it is also known as long form of the
periodic table.

Main Features of the Modern Periodic Table


Table 1.1
(i)It has groups and periods like Mendeleev’s
Periodic table.

Period
Nature
Elements
(ii)It has nine groups, I to VIII and zero group.
1
Shortest period 2 H and He
(Noble gases were known when the table was
2
Short period
8 (3)Li to (10)Ne
made)
3
Short
Period
8
Na to (18)Ar
(iii)Zero groups contains noble gases (He, Ne, Ar,
(11)
Kr, Xe and Rn)
4
Long Period
18 (19)K to (36 )Kr
(iv)Groups from I to VII are divided into sub5
Long Period
18 (37)Rb to (54)Xe
groups A and B.
6
Longest Period 32 (55)Cs to (86)Rn
(v) Zero group and group VIII has no sub-groups.
7

Incomplete
23 (87)Fr to (111)Rg
(vi) This table has seven periods.
(vii) Elements from atomic number 84 (Polonium)
onwards are radioactive. A few others having lower atomic number are also radioactive, TC (it is
synthetic also).
(viii) Elements after atomic number 92 (Uranium) are all synthetic or man made. They have been produced
by nuclear reactions.

IUPAC Table
It is simply a changed version of the long form of the periodic table. It was recommended in 1984 by IUPAC
committee.


www.pdfgrip.com
Periodic Table and Periodicity of Properties  3

Main Features
(i) This table has 18 groups, 1 to 18.
(ii) The notation IA – VIIA, IB – VIIB and VIII has been dropped.
(iii) Elements of group 1 and group 2 and 13 to 17 have all their inner shells completely filled. These are
s- and p- block elements.
(iv) Elements of group 18 are noble gases. They have all their orbitals completely filled.
(v) Elements belonging to group 3 to group 12 have their inner (n-1)d or (n-2)f orbital partially filled (these
are d- and f-block elements).

Periodic Table and Aufbau Principle
Electron configuration of elements and the modern periodic table (or
1s1
H

1s22s22p1 B
IUPAC table) correspond with each other. The electronic configuration
1s2
He 1s22s22p2 C
of atoms follows Aunfbau principle. It states that electrons should be
1s22s1 Li
1s22s22p3 N
placed, one at a time, in the lowest energy orbital. When an orbital is
1s22s2 Be 1s22s22p4 O
fully filled then the electron should occupy next higher energy orbital.
The sequence followed is 1s<2s<2p<3s<3p<4s<3d<4p<5s etc. Thus,
addition of each electron in an orbital makes an element.
Formation of other elements can be seen in the same way.
In the modern periodic table, the first period has 2 elements and it is due to gradual filling of 1s orbital,
1s1 and 1s2 i.e., H and He.
The second period contains 8 elements only due to gradual filling of 2s 2p orbital i.e., 2s1, 2s2, 2s22p1,
2
2s 2p2, …………2s22p6.

Note
•

•
•

Table 1.2

In 3rd period, there are 8 elements only due to
gradual filling of 3s3p orbitals. The 3d orbital has
higher energy than 4s and when 4s filled period is

changed. Thus period of an element corresponds to
highest orbit present in electron configuration of
the atom of that element.
Similarly fifth period has only 18 elements due to
filling of 5s 4d and 5p orbitals.
In the 6th period, orbitals filled are 6s2 4f14 5d1o and
6p6. So there are 32 elements.

Therefore, it may be said that the modern periodic table
is a consequence of Aufbau principles.

Period
1

Electron
Configuration
1s1 to 1s2

Number of
Elements
2

3

2s1 to 2s22p6
3s1 to 3s23p6

4

4s to 3d 4s 4p


18

5
6

5s to 4d 5s 5p
6s1 to 4f14 5d1o
6s2 6p6

18
32

2

1
1

1o
1o

8
8
2
2

6

6


Determination of group and period of an element
Group and period
Table 1.3
Modern Period Table

IUPAC Table

Group group = total valence electrons

s-block , group = total valence electron

Sub-group A = when valence orbital are s or s and p

p-block, group = total valence electron +10
(Continued)


www.pdfgrip.com
4  Basic Concepts of Inorganic Chemistry
Modern Period Table

IUPAC Table

Sub-group B = when valence orbitals are (n-1)d ns.
Period: Equals highest orbit number present in electron
configuration of atom of the element

d-block, group = total valence electrons.
Same as in modern periodic table.


Note:
•

The f-block elements are placed in group 3rd (modern and IUPAC table both)
Lanthanides in 6th period and actinides in 7th period.

Relationships in the Periodic Table
(1) Group relation
Elements of a group are similar.
Group1 
  Li Na K Rb Cs
Group2 
  Be Mg Ca Sr Ba etc., are similar. However, first elements of each main group
(i.e., s and p-block) differ from rest of the group members, Example, Li, Be, B, C, N, O, and F differ
from their higher group members respectively.
Li differ from Na, Be from Mg, etc. This difference is due to:
(a) Small size (of the first member)
(b) High electronegativity (of the first member)
(c) Non-availability of d orbitals in the valence shell of these elements i.e., 2nd period elements).
Valence orbitals are s and p only.
(2) Horizontal relationship
(a) Elements of Group VIII (or 8, 9 and 10 of IUPAC table) have horizonal similarity
  Fe  CO  Ni   Ru  Rh  Pd  Os  Ir  Pt



(b)Transition elements form a horizonal series.
Ti, V, Cr, Mn, Fe, Co, Ni, etc and are similar
(c) Lanthanides i.e., elements from Ce to Lu, are similar in many ways.
(3) Diagonal relationship

Some elements placed diagonally in the table are similar.
Table 1.4
Gr →
Elements

1(IA)

2 (IIA)

3 (IIIA)

4 (IVA)

Li

Be

B

C

Na

Mg

Al

Si

Li is similar to Mg, Be to Al and B to Si. It is known as diagonal relationship.

Diagonal similarity is due – either
(a) To similar ionic radii
Li+ = 0.76A, Mg2+ = 0.72Å
(b)or to similar ionic potential


www.pdfgrip.com
Periodic Table and Periodicity of Properties  5
Ionic pot. = Charge/ionic radius
Be2+ = 6.6, Al+3 = 6.0

Classification of Elements
Different concepts can be applied for classification of elements. Example, properties, electron configuration,
physical state of elements, atomicity etc.

(A) Properties and Type of Elements
Elements are of the following four types:
(I) Normal or Representative Elements
Elements belonging to groups IA to VIIA of the modern periodic table. (or 1, 2, and 13 to 17 of IUPAC
table) are called ‘Normal Elements’. Some of the groups have got special names.

Note:
(i) Alkali metals because oxides and hydroxides
are basic and water soluble.
(ii) Alkaline earth because oxides and
hydroxides are partially soluble like earth
(soil).
(iii) Pnicogens because some compounds have
very bad odour.
(iv) Chaleogens because they form ores.

(v) Halogens because they form salt like
compounds.

Table 1.5
Group

Name

Electronconfiguration
ns1

IA

Alkali metals

IIA

Alkaline earth metals

ns2

IIIA

No special name

ns2np1

IVA

No special name


ns2np2

VA

Pnicogens

ns2np3

VIA

Chalcogens

ns2np4

VIIA

Halogens

ns2np5

(II) Noble gases
Elements of zero group (or 18th group of IUPAC) are known as “Noble Gases”. They are also called
‘Rare Gases’ or ‘Inert Gases’. The gases are (2)He, (10)Ne, (18)Ar, (36)Kr, (54)Xe and (86)Rn. These elements
have their outermost orbitals full filled.
He 1s2, Ne to Rn – ns2np6.
(III) Transition Elements
Elements having partially filled (n-1) d
orbitals in their atoms or ions are known as
‘Transition Elements’.

General electron configuration (n–1)
dxns2 or (n–1)dxns1 [x = 1 to 10].
There are four such series in the table:
All these elements are metals. The only
member liquid is Hg, rest are solids. They
have many common characteristics (Cf Tr
elements).

Table 1.6
Series

Elements

1st series (3d)

21

2nd series (4d)

39

3rd series (5d)

57

4th series (6d)
[incomplete]

No. of
Elements


Sc to 30Zn

10

Y to 48Cd

10

La, 72Hf to 80Hg

10

Ac, 104Rf to 111Rg

7

89


www.pdfgrip.com
6  Basic Concepts of Inorganic Chemistry
(IV)  Inner Transition Elements
Elements which have incompletely filled (n – 2)f and (n – 1)d are called ‘Inner Transition Elements’.
General e-configuration – (n – 2) fx (n – 1)d1 ns2, x = 1 to 14
However, there is variation in the electron configuration of these elements (it will be discussed
in Lanthanides).
There are two such series in the table, corresponding to filling of 4f and 5f orbitals.
Lanthanides (4f)
(La) 58Ce to 71Lu

Actinides (5f)
(Ac) 90Th to 103Lw
Elements after
U(92) are synthetic and known as ‘Transuranic Elements’.

(B) Electron-Configuration and Type of Elements
(I)  s-block elements
When last electron of the configuration is added to s - orbital, the element belongs to s-block.
Gr 1 (ns1) Li Na K Rb Cs Fr and Gr 2 (ns2) Be Mg Ca Sr Ba Ra elements are s-block elements.
These elements are most electropositive and highly reactive.
(II)  p-block elements
In these elements last electron is added to a p-orbital. This block has variety of elements.
Table 1.7
ns2np1

ns2np2

ns2np3

ns2np4

ns2np5

ns2np6

IIIA (13)

IVA (14)

VA (15)


VIA (16)

VIIA (17)

Zero (18)
He

(a)
(b)
(c)
(d)

B

C

N

O

F

Ne

Al

Si

P


S

Cl

Ar

Ga

Ge

As

Se

Br

Kr

In

Sn

Sb

Te

I

Xe


Tl

Pb

Bi

Po

At

Rn

Elements below the line joining B to At are metals, (Al to Tl; Sn and Pb and Bi; and Po).
Elements above the line joining B to At are non-metals (B, C, N., O, F, P, S, Cl, Br, I).
Elements falling on the line or very near to it are metalloids (Si, Ge, As, Sb, Se, Te).
Noble gases are also taken as non-metals.

(III)  d-block elements
The last electron in the configuration is filled in (n–1) d orbitals.
General configuration – (n – 1)dxns2 or (n – 1)dxns1 [x = 1 to 10].
There are four such series in the periodic talbe (3d, 4d, 5d and 6d); i.e., transition elements are
d-block elements. These elements have common properties (details in the chapter general properties of
transition elements).


21

44.96


39

40.08

38

39.10

37

88

226.03

Ra

Actinoids

88.91

Y

Ac–Lr

La–Lu

Lanthanoids

223


Fr

87

132.91 137.34

Ba

56

Cs

87.62

55

Sr

85.47

Rb

Sc

20

Ca

19


K

3

24.31

22.99

Mg
6

7

8

9

10

11

12

105

104

W

74


95.94

42

Mo

52.01

24

Cr

Re

75

98.91

43

Tc

54.94

25

Mn
45


Rh

58.93

27

Co
46

Pd

58.69

28

Ni
47

Ag

63.54

29

Cu
48

Cd

65.41


30

Zn
In

49

69.72

31

Ga

50

Sn

72.59

32

Ge

51

Sb

74.92


33

As

30.97

P

91

90

60

Nd

[264]

Bh

107

I

53

79.91

35


Br

35.45

Cl

17

19.00

9

F

17

54

Xe

83.80

36

Kr

39.95

18


Ar

20.18

10

Ne

4.00

76

Os

Ir

77

78

Pt

79

Au

80

Hg


Tl

81

82

Pb

83

Bi

61

Pm

[277]

Hs

108

62

Sm

[268]

Mt


109

63

Eu

[271]

Ds

110

64

Gd

[272]

Rg

111

65

Tb

[285]

112


Uub

66

Dy

67

Ho

68

Er

69

Tm

210

84

Po

70

Yb

210


85

At

71

Lu

222

86

Rn

Th

Pa

U

92

93

Np

94

Pu


95

Am

96

Cm

97

Bk

98

Cf

99

Es

100

Fm

101

Md

227.03 232.04 231.04 238.03 237.05 239.05 241.06 244.07 249.08 252.08 252.09 257.10 258.10


Ac

259

No

102

Lr

262

103

138.91 140.12 140.91 144.24 146.92 150.36 151.96 157.25 158.92 162.50 164.93 167.26 168.93 173.04 174.97
89

59

Pr

58

Ce

57

[266]

Sg


106

La

[262]

Db

[261]

Rf

52

Te

78.96

34

Se

32.06

S

16

16.00


8

O

16

2

He

18

101.07 102.91 106.42 107.87 112.40 114.82 118.71 121.75 127.60 126.90 131.30

44

Ru

55.85

26

Fe

28.09

Si

15


14.01

7

N

15

178.49 180.95 183.85 186.21 190.23 192.22 195.08 196.97 200.59 204.37 207.19 208.98

Ta

73

Hf

92.91

72

41

Nb

50.94

V

23


91.22

40

Zr

47.90

Ti

22

26.98

Al

14

Na

12.01

13

6

C

14


10.81

5

B

13

12

5

Relative atomic mass, Ar

9.01

4

1.008

Element symbol

Atomic number, Z

11

Be

4


2

H

1

6.94

Li

3

1.008

H

1

1

Periodic table

www.pdfgrip.com


www.pdfgrip.com
8  Basic Concepts of Inorganic Chemistry
(IV)  f–block elements
The last electron in the configuration is added to (n – 2)f orbitals.

General configuration – (n – 1)f x (n – 1)d1ns2 [x = 1 to 14].
Lanthanides and Actinides are f-block elements. All elements are metals. All elements of actinide
series are radioactive. Elements after U (92) are synthetic. These are gr IIIB (3 of IUPAC) elements of
the table. However, to keep the symmetry of the table scientific, the elements are placed at the bottom
of the table.

Atomicity and Elements of the Periodic Table
Monoatomic Elements
Noble gases He, Ne, Ar, Kr, Xe and Rn, are monoatomic in their standard state. It is because noble gases have
closed shell electron configuration (1s2 or ns2np6), and are not in a position to form bond with each other.
However, in excited states species like He+2 can exist, such species are known as “Exonomers”. In the vapour
state, Hg (5d1o 6s2) is also monoatomic.

Diatomic molecules: (H2, D2, HD, HT, DT, T2, N2, O2, F2, C2, Br2, I2).
Diatomicity is also related with electron configuration. Hydrogen and halogens achieve stable electron –
configuration by forming a single electron – pair bond in a diatomic molecule. For Nitrogen (2s22p3) and
Oxygen (2s2 2p4), multiple bonding gives stable e – configuration, so diatomicity, i.e., H2, X2, N2, O2.
Phosphorus and Sulphur also form P2 and S2 at high temperature but no at 25oC.
Discrete Polyatomic molecules: (P4, S8, Se8 etc.)
Dinitrogen (N2) and dioxygen (O2) are stable due to very effective π bounding involving p – orbitals i.e.,
(p – p)π bonding. For the third period elements (p – p)π bonding is not effective due to larger size of p –
orbitals of P, S etc., and more core electrons (8 es).
Therefore, these elements, instead, form discrete polyatomic
3
molecules like P4, S8. White phosphorus P4 is tetrahedral. Each ‘P’ forms
three bonds, with the P – P distance 2.21Å and P – P – P angles 60o.
3
3
This small angle is associated with considerable strain in the ring
structure. Therefore, white phosphours is highly reactive. Arsenic and

antimony are As4 and Sb4.
3
The most common form of S is S8. The S8 has crown ring structure,
(given below) such rings are stacked over each other in the solid structure
Figure 1.1
S is sp3 in the S8 ring and the angle SSS is around 108o. Selenium is
Se8 like S8.
Carbon also form polyatomic molecules, C60 called “Fullerenes”. The C60 has soccer ball structure
(details in allotropy).
Gaint molecules
Elements which can form 2, 3, 4 etc covalent bonds can also form gaint
(or macromolecular) species. Thus p – block elements have this property
like B, C, P, S etc. have this properties.
B is inert and has very high m. p. It is due to the fact that boron
exists as B12, icosahedral structure.
Other macromolecular species include
(i) Diamond,
(ii) Graphite,

Figure 1.2


www.pdfgrip.com
Periodic Table and Periodicity of Properties  9
( iii) SiC,
(iv) Black phosphorus,
(v) Catena sulphur, chain form of sulphur. It is present in plastic sulphur.

Periodicity in the Periodic Table
Effective Nuclear Charge

In a multielctronic atom electrons of the inner shells put a screen of negative charge on the nucleus. Therefore,
outer electrons experience lesser attractive effect of nucleus. The actual nuclear charge an outer electron
experiences is called ‘Effective Nuclear Charge’ (Z*).
It is given as, Z* = Z – σ, where Z = atomic number, σ = screening (or shielding) constant.
Thus, if σ is known, Z* can be calculated. It is calculated using Stater’s rule. The rules are:
(i) Orbitals are grouped in the following way
(1s) (2s,2p) (3s3p) (3d) (4s4p) …………. etc.
(ii) σ = 0, for e in question (i.e., e of nth orbit) or electrons in the higher group (i.e., n + 1)
(iii) σ = 0.35 per electron for rest of the group electrons.
(iv) σ = 0.3, for a 1s electron screening another 1s electron.
(v) σ = 0.85 per electron for all electrons in (n – 1)th group.
(vi) σ = 1, per electron for all electrons in (n – 2) or lower groups.
(vii) In case of electron being shielded of nd or nf, all electrons lying to the left of nd or nf gr, σ = 1 per
electron.
The above rules are summerized in the table below:
Table 1.8
Electron
group

All higher
groups

Same
Group
Group with
group with (n – 1)
(n – 2)

1s


0

0.30

0

0

ns, np

0

0.35

0.85

1.00

nd, nf

0

0.35

1.00

1.00

The screening power of different orbitals follows the order: s > p > d > f
Examples

(a) Z* for valence electron of Na

Na (1s2) (2s2 sp6) (3s1)
11

σ = 1x2 + 0.85x8 + 0.35x0
= 2 + 6.8 + 0
= 8.8

Z* = 11.0 – 8.8 = 2.2
(b) Z* for k, valence electron

k (1s2) (2s22p6) (3s23p6) (4s1)
19

σ = 1x10 + 0.85x8 + 0.35x0
= 10 + 6.8 + 0
= 16.8

Z* = 19.0 – 16.8 = 2.2

(c) Li Z*

1s2 2s1

σ = 0.85x2 + 0.35x0

= 1.70 + 0

= 1.7


Z* = 3.0 – 1.7 = 1.3
(d) Z* Be

1s2 (2s2)

σ = 0.85x2 + 0.35x1
= 2.05

Z* = 4.00 – 2.05
= 1.95
(e) Z*Mg

1s2(2s2 2p2) (3s2)


www.pdfgrip.com
10  Basic Concepts of Inorganic Chemistry

σ = 1x2 + 0.85x8 + 0.35x1
= 2 + 6.8 +0.35
= 9.15

Z* = 12.00 – 9.15
= 2.85

(f) Z* for 10th d – electron of Zn(30)

Zn= 1s2 (2s22p6) (3s23p6) 3d1o 4s2
30

  σ = 1x18 + 0.35x9 + 2x0

= 18 + 3.15

= 21.15
  Z* = 30.00 – 21.15

= 8.85

Variation of Z* in the Periodic Table
(i) Z* remains constant in a group (except the first member)
Li
Na
K
Rb
Cs

Z* 1.3
2.2
2.2
2.2
2.2
(ii) Z* increases in a period from left to right

Li Be
B
C
N
O
F


Z*→ 1.3 1.95
2.6
3.25
3.9
4.55
5.2

Atomic and Ionic Size
Wave mechanical model of atom shows that a precise boundary around an atom can not be drawn, because
electron probability distribution never becomes exactly zero. Therefore, radius of an atom is difficult to
determine. Radii of atoms are determined in different combined states and so defined as:
(i) Covalent radius
(ii) Van der Waal’s radius
(iii) Crystal radius
(i)  Covalent Radius
Half of the bond length between two similar covalently
bonded atoms is defined as covalent radius. It is based on
the assumption that atoms are spherical The Cl–Cl bond
length is found 1.988 Å, so covalent radius of chlorine is
0.99 Å.
Thus , rcovalent = Bond length/2
Bond length = 1.988 Å
c
Covalent radius = 1.988/2 = 0.99Å
ERQGOHQJWK

In a similar way radii of other atoms are determined
Covalent radii are commonly referred to as atomic
Figure 1.3

radii.
Elements may have multiple bonds and they are always shorter than the corresponding single
bonds,
C – C (1.54Å), C = C (1.33Å), C ≡ C (1.2Å)
N – N (1.45Å), N = N (1.25Å), N ≡ N (1.10Å)
Therefore, double and triple bond radii are also defined (half the double bond length and triple
bond length), e.g.,
Double bond radius of C = 1.33/2 = 0.665Å
Triple bond radius of N = 1.10/2 = 0.55Å
The double and triple bond radii of an atom are approximately 0.87 and 0.78 times the single bond
radii, respectively.


Ce
1.65

Lanthanide

3
Sc
1.44
Y
1.62
La
1.68
Ac

Be
1.25
Mg

1.45
Ca
1.74
Sr
1.91
Ba
1.98

Ra

Fr

Group
1
H
0.37
Li
1.34
Na
1.57
K
2.03
Rb
2.16
Cs
2.35

Pr
1.64


4
Ti
1.32
Zr
1.45
Hf
1.44

Nd
1.64

5
V
1.22
Nb
1.34
Ta
1.34

Pm
-

6
Cr
1.17
Mo
1.29
W
1.30


Sm
1.66

Eu
1.85

Group
7
8
Mn
Fe
1.17 1.17
Te
Ru
1.24
Re
Os
1.28 1.26

*Radii of group 18 elements is Van der Walls’ radii.

7

6

5

4

3


2

1

Period

Covalent radii of the elements (Å)

Gd
1.61

9
Co
1.16
Rh
1.25
Ir
1.26

Tb
1.59

10
Ni
1.15
Pd
1.28
Pt
1.29


Dy
1.59

11
Cu
1.17
Ag
1.34
Au
1.34

Table of Radii

Ho
1.58

12
Zn
1.25
Cd
1.41
Hg
1.44

Er
1.57

B
0.90

Al
1.25
Ga
1.20
In
1.50
Tl
1.55

Group
13

Tm
1.56

C
0.77
Si
1.17
Ge
1.22
Sn
1.40
Pb
1.46

Group
14

Yb

1.70

N
0.74
P
1.10
As
1.21
Sb
1.41
Bi
1.52

Group
15

Lu
1.56

O
0.74
S
1.04
Se
1.14
Te
1.37
Po
-


Group
16

F
0.72
Cl
0.99
Br
1.14
I
1.33
At
-

Group
17

Rn
-

Xe
2.18

Kr
1.98

He
1.4
Ne
1.54

Ar
1.92

Group
18

www.pdfgrip.com
Periodic Table and Periodicity of Properties  11