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The Pearson Guide to

Inorganic Chemistry for
the Medical Entrance
Examinations

Atul Singhal


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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
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ISBN 9788131787830
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Contents
Chapter 1: Chemical Bonding

1.1–1.57

Chapter 2: Classification of Elements and Periodicity Properties

2.1–1.32

Chapter 3: Hydrogen and Its Compounds

3.1–3.34

Chapter 4: S-Block Elements Group I

4 .1–4.63

Chapter 5: Boron Family IIIA – Group Elements

5.1–5.35

Chapter 6: Carbon Family IVA – Group Elements

6.1–6.54

Chapter 7: Nitrogen Family VA – group Elements

7.1–7.51


Chapter 8: Oxygen Family VIA – group elements

8.1–8.42

Chapter 9: Halogen Family VIIA – group elements

9.1–9.41

Chapter 10: Noble Gases or Zero Group VIIIA – Group elements

10.1–10.21

Chapter 11: Transition Metals Including Lanhanides and Actinides

11.1–11.40

Chapter 12: Co-ordination Chemistry

12.1–12.53

Chapter 13: Chemistry of Heavier Elements

13.1–13.47

Chapter 14: Principles of Qualitative Analysis

14.1–14.23


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Preface
The Pearson Guide to Inorganic Chemistry for the Medical Entrance Examination is an invaluable
book for all the students preparing for the prestigious medical entrance examination. It provides
class-tested course material and problems that will supplement any kind of coaching or resource the
students might be using. Because of its comprehensive and in-depth approach, it will be especially
helpful for those students who do not have enough time or money to take classroom courses.

•

A careful scrutiny of previous years’ A.I.P.M.T papers and various other competitive PMT
examinations during the last 10 to 12 years was made before writing this book. It is strictly
based on the latest AIPMT/STATE P.M.T syllabus (2009–10) recommended by the executive
board. It covers the subject in a structured way and familiarizes students with the trends in these
examinations. Not many books in the market can stand up to this material when it comes to the
strict alignment with the prescribed syllabus.

•

It is written in a lucid manner to assist students to understand the concepts without the help of
any guide.

•

The objective of this book is to provide this vast subject in a structured and useful manner so as

to familiarize the candidates taking the current examinations with the current trends and types
of multiple-choice questions asked.

•

The multiple-choice questions have been arranged in following categories:
3 Gear Up I (To Revise the Concepts)
3 Gear Up II (To Sharpen the Concepts)
3 Gear Up III (Concept Crackers)
3 A Peep into the AIPMT
3 MCQ’s from Recent Entrance exams
3 Assertion and Reason

This book is written to pass on to another generation, my fascination with descriptive physical
organic chemistry. Thus, the comments of the readers, both students and instructors, will be sincerely
appreciated. Any suggestions for added or updated additional readings would also be welcome.

Dr Atul Singhal




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Acknowledgements
The contentment and ecstasy that accompany the successful completion of any work would remain
essentially incomplete if I fail to mention the people whose constant guidance and support has
encouraged me.
I am grateful to all my reverend teachers, especially, the late J. K. Mishra, Dr D. K. Rastogi, the late
A. K. Rastogi and my honourable guide, Dr S. K. Agarwala. Their knowledge and wisdom has continued to assist me to present this work.

I am thankful to my colleagues and friends, Deepak Bhatia, Er Vikas Kaushik, Er A. R. Khan, Vipul
Agarwal, Er Ankit Arora, Er Wasim, Manoj Singhal, Vijay Arora (Director, Dronacharya), Anupam
Shrivastav (Bansal Classes), Rajiv Jain (MVN, Faridabad), Ajay Verma, Ashutosh Tripathi, Vivek
shukla and Gaurav Bansal (C-25Classes) Satish Gupta (Resorance Jaipur), Chandan Kumar (Everonn
Toppers).
I am indebted to my father, B. K. Singhal, mother Usha Singhal, brothers, Amit Singhal and Katar
Singh and Sisters, Ambika and Poonam, who have been my motivation at every step. Their neverending affection has provided me with moral support and encouragement while writing this book.
Last but not the least, I wish to express my deepest gratitude to my wife Urmila and my little,—but
witty beyond years, daughters, Khushi and Shanvi who always supported me during my work.

Dr Atul Singhal




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C HAPTER

1

CHEMICAL BONDING
CHAPTER CONTENTS
3 Kossel-Lewis approach to chemical bond formation, ionic bonds, covalent bonds,
polarity of bonds and concept of electronegativity, valence shell electron pair
repulsion (VSEPR) theory, shapes of simple molecules, valence bond theory,
hybridization involving s, p and d orbitals and shapes of molecules T and Q
bonds
3 Molecular orbital theory involving homounclear diatomic molecules
3 Hydrogen-bonding


BONDS
VALENCY
Valency is a property of atoms whereby they form chemical
bond among themselves. The term valency was introduced
by Frankland and it means ‘power to combine.’ Hence, it is
the power of an atom to combine with another atom. Atoms
do so by either giving up or accepting electrons in their
outermost shell. Modern or electronic concept of valency
was given by Kossel and Lewis; it was completed by
Langmuir.
Valency (V)  No. of valence electrons
For instance, the electronic configuration for the group
IA element sodium (Na), is 2, 8, 1. Here, the number of
valence electron is 1 and hence its valency is 1.
If the number of valence electrons is more than 4, then
we use the following relationship to determine the valency:


V  V e  8 (number of valence electrons minus 8)

For example, the configuration of nitrogen (N) is 2, 5. According to the above relationship, its valency will be given as
V  5  8  3 (negative sign signifies the tendency
to accept electrons)

CHEMICAL BOND
Chemical bond is the force of attraction that binds two
atoms together. A chemical bond balances the force of
attraction and force of repulsion at a particular distance.
A chemical bond is formed to:

• Kattain the octet state
• minimize energy
• gain stability
• decrease reactivity
When two atoms come close to each other, forces of
attraction and repulsion operate between them. The dis-


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1. 2

Chemical Bonding

tance at which the attractive forces overcome repulsive
forces is called bond distance. Here, potential energy for
the system is lowest, hence the bond is formed.

Mg

(2, 8, 2) (2, 6)
Al

Types of Bonds
Following are the six types of chemical bonds. Here, they
are listed in a decreasing order of their respective bond
strengths.
1.
2.
3.

4.

Octet Rule
It was introduced by Lewis and Kossel. According to
this rule, each atom tries to obtain the octet state, that is,
a state with eight valence electrons.

Exceptions to the octet rule

•
•

Transition metal ions like Cr3, Mn2, Fe2.
Pseudo inert gas configuration cations like Zn2+, Cd2.

Contraction of octet state

•

Here central atom is electron deficient or does not
have an octet state. For example,
BeX 2 BX 3 AIX 3 Ge(CH 3 )3
4
6
6
6 e

Expansion of octet state

•

•
•
•

Here central atom has more than 8 e due to empty d
orbital. For example, PCl5, SF6, OsF8, ICl3
Odd electronic species like NO, NO2, ClO2
Inter halogens compounds like IF7, BrF3
Compounds of xenon such as XeF2, XeF4, XeF6

IONIC OR KERNEL BOND
Ionic bond is formed by the complete transfer of valence
electrons from a metal to a non-metal. This was first studied by Kossel.
For example,
Na

Cl
(2, 8, 1) (2, 8, 7)

•

•

5. Hydrogen bond
6. Van der Waals bond
Metallic bond, hydrogen bond and van der Waals bond
are interactions.

Na Cl
(2, 8) (2, 8)


(2, 8) (2, 8)
Al3 N3

 N

(2, 8, 3) (2, 5)

•

Ionic bond
Covalent bond
Coordinate bond
Metallic bond

Mg2 O2

 O

•
•

(2, 8) (2, 8)

Number of electrons transferred is equal to electrovalency.
Maximum number of electrons transferred by a metal
to non-metal is three, as is the case of AlF3, (Al metal
transfers three electrons to F).
During electron transfer, the outermost orbit of metal
is destroyed and the remaining portion is called core or

kernel, so this bond is also called kernel bond.
Nature of ionic bond is electrostatic or coloumbic force
of attraction.
It is a non-directional bond.

Conditions for the Formation of an Ionic Bond
The process of bond formation must be exothermic
($H  ve) and for it the essential conditions are

•
•
•
•
•

Metal must have low ionization energy.
Non-metal must have high electron affinity.
Ions must have high lattice energy.
Cation should be large with low electronegativity.
Anion must be small with high electronegativity.

Born–Haber Cycle
The formation of an ionic compound in terms of energy
can be shown by Born–Haber cycle. It is also used to find
lattice energy, ionization energy and electron affinity.
For example,
Ionization

M(s) }Sublimation
}}}}

m M(g) } }}}m M(g) + e
+1

+s

decomposition

Addition of e X(g)
/2 X2 } }}}}}
m X(g) } }}}}}
m
+E

1

+s

Crystal formation

M(g) 1 X(g) } }}}}}}
m MX(g)
U

$Hf  S  1/2 D  I  E  U
Here,
S  Heat of sublimation
D  Heat of dissociation
I  Ionization enthalpy
E  Electron gain enthalpy or electron affinity
U  Lattice energy



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1. 3

Chemical Bonding

•

For the formation of an ionic solid, energy must be
released during its formation, that is, %H must be negative for it.
E  U > S
1/2 D
I

NaCl
MgCl2
AlCl3
SiCl4
As charge on a metal atom increases, its size decreases.
In case of univalent and bivalent ionic compounds, lattice energy decreases as follows:
Bi-bi  Uni-bi or Bi-uni  Uni-uni

Properties of Ionic Compounds
1. Ionic compounds have solid crystalline structures (flat
surfaces), with definite geometry, due to strong electrostatic force of attraction as constituents are arranged in
a definite pattern.
2. These compounds are hard in nature.
Hardness t Electrostatic force of attraction








t

Charge on ion












t


1


Ionic radius

3. Ionic compounds have high value of boiling point, melting point and density due to strong electrostatic force of
attraction.
Boiling point, melting point о Electrostatic force of
attraction
1
Volatile nature о
Electorstatic force of attraction
4. Ionic compounds show isomorphism, that is, they have
same crystalline structure. For example, all alums, NaF
and MgO.
5. These are conductors in fused, molten or aqueous state
due to presence of free ions. In solid state, these are nonconductors as no free ions are present.

For example,

MgO  MgCl2  NaCl.
9. Ionic compounds are soluble in polar solvents like water
due to the high dielectric constant of these solvents,
therefore, force of attraction between ions are destroyed
and they dissolve in the solvent.
Facts Related to Solubility
• If %H (hydration)  Lattice energy then ionic compound is soluble.

•

If %H (hydration)  Lattice energy then ionic compound is insoluble

•

If %H (hydration)  Lattice energy then the compound is at equilibrium state

Some Solubility Orders
a.

LiX  NaX  KX  RbX  CsX

b.

LiOH  NaOH  KOH  RbOH  CsOH

c.

BeX2  MgX2  CaX2  BaX2

d.


Be(OH)2  Mg(OH)2  Ca(OH)2  Ba(OH)2

e.

BeSO4  MgSO4  CaSO4  SrSO4  BaSO4

f.

AIF3  AlCl3  AlBr3  AII3

6. They show fast ionic reactions as activation energy is
zero for ions.
7. They do not show space isomerism due to nondirectional nature of ionic bond.
8. Lattice energy (U) is released during the formation of an
ionic solid molecule from its constituent ions.

•

Lattice Energy
Lattice energy is also the energy needed to break an
ionic solid molecule into its constitutent ions. It is
denoted by U.

•

Crystals of high ionic charges are less soluble. For
example, compounds of CO32, SO42, PO43 are less
soluble.
Compounds Ba2, Pb2 are insoluble as lattice

energy  %Hhy
Compounds of Ag (salt) are insoluble as lattice
energy  %Hhy
Presence of common ions decrease solubility. For
example, solubility of AgCl decreases in presence of
AgNO3 or KCl, due to presence of common ions that
is, Ag and Cl respectively.

U о Charge on ion
 о

1
Size of ion

Hence, lattice energy for the following compounds
increases in the order shown below:

COVALENT BOND
A covalent bond is formed by equal sharing of electrons
between two similar or different atoms.


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1. 4

Chemical Bonding

•


If atoms are same or their electronegativity is same, the
covalent bond between them is non-polar. For example,

HH (too close)
+

X–X, O = O, N y N
If atoms are different or have different value of electronegativity, the covalent bond formed between them is
polar. For example,

E

E


E

HOH,

•
•
•
•
•
•


E

E


HX

Here, number of electrons shared or covalent bonds represent covalency.
One atom can share a maximum of three electrons with
the other atom.
The nature of covalent bond is explained on the basis of
Heitler–London’s valence bond theory, Pauling–Slater’s
overlapping theory and Hund–Mullikan’s theory.
Orbital concept of covalent bond was introduced by
Heitler and London. According to this concept, “Covalent bond is formed due to half-filled atomic orbitals
having electrons with opposite spin to each other.”
Due to overlapping, the potential energy of system
decreases.
The internuclear distance with maximum overlapping
and greater decrease of potential energy is known as
bond length.

Energy consideration of covalent bond When two
hydrogen atoms HA and HB with respective electrons eA and
eB approach each other, following attractive and repulsive
force start operatings.
Attractive

Electron cloud

Repulsive
Nucleus
HA


HB

Figure 1.1
Here, attractive forces between HAeA and HBeB and HBeA
and HAeB.
Repulsive forces are between eA and eB and between
nucleus of HA and HB.
It is observed that attractive force are more than a repulsive forces which results in decreased energy, so the potential energy of the system decreases.
The minimum energy point corresponds to critical distance between two nuclei, when maximum lowering of
energy takes place. This distance is called bond length
e.g., in H–H, bond length is 74 pm.

H

H (too far)

0

Energy

•

H H
–
Bond length Inter nuclear
(74pm) distance

Figure 1.2

Features of Covalent Compounds

1. Covalent compounds mostly occur in liquid and gaseous
state, but if molecular weight of the compound is high,
they may occur in solid state also. For example,
Molecular wt
Glucose
Sugar
180
342
(less solid)
(more solid)
2. ‘Like dissolves like’, that is, non-polar solute dissolves
in non-polar solvent. For example, CCl4 dissolves in
organic solvents.
Similarly, polar solutes dissolves in polar solvent. For
example, alcohol and ammonia dissolve in water.
3. Covalent compounds have lower boiling point and melting point values than those of ionic compounds. This is
because covalent bond is weak van der Waals force in
nature.
Comparable to
KOH

HX
Strong
Weak
ionic force
van der
of attarction
Waals forces
Boiling point and melting point о Hydrogen bonding
о Molecular weight

For example,
HF 
HI

HBr

HCl
Due to
As molecular
H-bonding
weight decreases
4. Covalent compounds are non-conductors due to absence
of free ions, but graphite is a conductor, as in graphite,
free electrons are available in its hexagonal sheet like
structure.
In case of diamond, the structure is tetrahedral so free
electrons are not available. It is therefore not a conductor.
5. Covalent bond is directional, so these compounds can
show space isomerism.


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1. 5

Chemical Bonding

6. When cation and anion are close to each other, the shape
of anion is distorted by the cation, this is called polarization. Due to this, covalent nature develops in an ionic
molecule.

Polarization t Covalent nature t


+

–

Polarization

1
Ionic nature

+

–

Distorted anoin

Figure 1.3 Effect of Polarization

Fajan’s rule
Polarization or covalent nature is explained by the following rules:
Charge on cation polarization, covalent nature or
polarizing power of cation t charge on cation. That is,
greater the charge on cation, greater will be its polarizing power and more will be covalent nature. For example, SiCl4  AICl3  MgCl2  NaCl
Size of Cation When charge is same and anion is com1
mon, consider it covalent nature t



Size of cation
That is, smaller cation has more polarizing power.
For example,

LiCl  NaCl  KCl  RbCl  CsCl
Max. covalent
Least ionic

Max. ionic
Least covalent

Li
 Na
 K
 Rb
 Cs

Smallest
in size

Largest
in size

Size of anion This property is taken into account when
the charges are same and the cation is common.
• Polarization or covalent nature t size of anion.
Hence, larger anions are more polarized.
For example, LiF