Chemistry

C. Calculate the oxidation number of 4. Balance the reactions/equations :
underlined atoms.
A. Balance the following reactions by
a. H2SO4 b. HH2N−S− O4O3 6 c. H3 2P−OO723- oxidation number method
e. f. C−r
d. NKa2 C−H22OP−4O4 a. Cr2O72 (aq) + SO32 (aq) Cr3⊕ (aq)
g.
+ SO42 (aq) (acidic)
D. Justify that the following reactions are
b. MnO4 (aq) + Br (aq) MnO2 (s)
redox reaction; identify the species + BrO3 (aq) (basic)

oxidized/reduced, which acts as an c. H2SO4 (aq) + C (s) CO2 (g) +

oxidant and which act as a reductant. SO2 (g) + H2O (l) (acidic)

a. 2Cu2O (s) + Cu2S (s) 6Cu (s) + SO2 (g) d. Bi (OH)3 (g) + Sn(OH)3 (aq) Bi (s)

b. HF (aq) + OH (aq) H2O(l) + F (aq) + Sn (OH)62 (aq) (basic)

c. I2 (aq) + 2 S2O32 (aq) S4O62 (aq) + 2I (aq) B. Balance the following redox equation by

E. What is oxidation? Which one of the half reaction method

following pairs of species is in its a. H2C2O4 (aq) + MnO4 (aq)
CO2(g) + Mn2⊕ (aq) (acidic)
oxidized state ?
b. Bi (OH)3 (s) +SnO22 (aq) SnO32 (aq)
a. Mg / Mg2⊕ b. Cu / Cu2⊕ + Bi (s) (basic)

c. O2 / O2 d. Cl2 / Cl

F. Justify the following reaction as redox 5. Complete the following table :

reaction. Assign oxidation number to the underlined
species and write Stock notation of compound
2 Na(s) + S(s) Na2 S(s)

Find out the oxidizing and reducing Compound Oxidation Stock
number notation
agents. AuCl3
SnCl2
G. Provide the stock notation for the V2O74
following compounds : HAuCl4, Tl2O, PtCl62
FeO, Fe2O3, MnO and CuO. H3AsO3

H. Assign oxidation number to each atom
in the following species.

a. Cr(OH)4 b. Na2S2O3 Activity :
c. H3BO3
I. Which of the following redox couple is Perform redox reaction experiment
with the help of Daniel cell under teacher
stronger oxidizing agent ? guidance.

a. Cl2 (E0 = 1.36 V) and Br2 (E0 = 1.09 V)
b. MnO4 (E0 = 1.51 V) and

Cr2O72 (E0 = 1.33 V)
J. Which of the following redox couple is
stronger reducing agent ?

a. Li (E0 = - 3.05 V) and

Mg (E0 = - 2.36 V)

b. Zn (E0 = - 0.76 V) and

Fe (E0 = - 0.44 V)

92

7. Modern Periodic Table

Can you recall? Mendeleev arranged the elements
known at that time in an increasing order of
• What was the basis of classification their atomic masses. The serial or ordinal
of elements before the knowledge of number of an element in the increasing order
electronic structure of atom? of atomic mass was referred to as its atomic
number. He folded this list in accordance
• Name the scientists who made the with recurrence of properties of elements and
classification of elements in the formed his periodic table consisting of vertical
nineteenth century. groups and horizontal series (now called
periods).
• What is Mendeleev’s periodic law? In Mendeleev’s periodic table, elements
• How many elements are discovered belonging to the same group showed similar
properties. Properties of elements in a series/
uptil now? period showed gradual variation from left to
• How many horizontal rows and vertical right. Mendeleev left some gaps corresponding
to certain atomic numbers in the periodic
columns are present in modern periodic table so as to maintain the periodicity of the
table? properties. Mendeleev’s periodic table was
7.1 Introduction accepted by the scientific community since
In the early nineteenth century about the newly discovered elements fitted well into
30 elements were known and were classified the gaps with their properties as predicted by
into three types on the basis of their physical Mendeleev’s periodic law. Inert gases, not
properties as: metals, nonmetals and predicted by Mendeleev and discovered in
metalloids. Subsequent noteworthy attempts later years also could be accommodated in this
For classification of the increasing number periodic table by creating an additional group.
of elements based on atomic mass were After the discovery of atomic structure,
Dobereiner’s triads and Newlands’ octaves. the atomic number, which was an ordinal
number assigned to element in Mendeleev’s
Just think periodic table, was recognized as the proton
number, Z, of that element. This was the
• How many days pass between two outcome of Henry Moseley’s work (1913) on
successive full moon nights? x-ray spectroscopic study of a large number of
elements. Moseley showed that the frequency
• What type of motion does a pendulum of x-ray emitted by an element is related to
exhibit? atomic number, Z, rather than the atomic mass.
The atomic number, Z, was considered as more
• Give some other examples of periodic fundamental property of the atom than the
events. atomic mass. As a result, Mendeleev’s periodic
law was modified. It is called the modern
In 1869, Russian chemist Dmitri periodic law and is stated as: “The physical
Mendeleev put forth periodic table of the 63 and chemical properties of elements are a
elements known at that time using the atomic
mass and properties of elements. Mendeleev’s periodic function of their atomic numbers”.
periodic table was based on the as Mendeleev’s
periodic law which is stated as “The physical
and chemical properties of elements are
periodic function of their atomic masses”.

93

Mendeleev’s periodic table was revised This numbering of the periods and groups
in accordance with the modern periodic law.
We will look into the final revised version is recommended by the International Union
of Pure and Applied Chemistry, IUPAC. The
known as the modern periodic table also boxes formed at the intersection of the periods
and groups are the places for individual
called the long form of periodic table in the
following sections. elements. Below the main table are placed two

7.2 Structure of the Modern Periodic Table series containing fourteen elements each.

Over a long period of time many There are in all 118 boxes to accommodate
scientists have come up with different forms
118 elements in the modern periodic table. As
of periodic table. However, the so called ‘long on today all the 118 boxes are filled as a result
form of periodic table’ or ‘the modern of discovery of manmade elements. IUPAC
periodic table’, which is a revised version has approved names and symbols of all the
of Mendeleev’s periodic table, is the most 118 elements. (Refer to Fig. 7.1)
convenient and widely used form of the The overall shape of the modern
periodic table of elements today. periodic table shows that it is divided into four
The modern periodic table has
horizontal rows intersecting the vertical blocks. Two groups on left form the s-block,
columns giving rise to a number of boxes.
The horizontal rows are called periods (which six groups on the right constitute the p-block,
Mendeleev called series) and the vertical
ten groups in the center form the d-block and
columns are called groups. There are seven the two series at the bottom constitute the

periods (numbered 1 to 7) and eighteen f-block (Fig. 7.2).
groups (numbered 1 to 18) in the modern
periodic table.

Periods Gro1ups 18
1 13 14 15 16 17
3 4
2 5 6 7 8 9 10 11 12
2

3

4
5

6

7

Series

Fig. 7.1 : Modern Periodic Table (Refer page 271)

94

7.3 Periodic Table and Electronic Along a period the atomic number
configuration increases by one and one electron is added
to outermost shell which forms neutral atom
Can you recall? of the next element. Every period ends with
complete octet configuration (or duplet in the
• What do the principal quantum number case of the first period) of the valence shell and
the next period begins with addition of electron
‘n’ and azimuthal quantum n u m b e r ‘ l ’ to the next shell of higher energy compared to
of an electron belonging to an atom the previous period. The first shell, thus, gets
filled along the first period. As the first shell
represent? can accommodate only two electrons, there
are two elements in the first period, namely, H
• Which principle is followed in the (Z=1) : 1s1 and He (Z=2) : 1s2. The first period
ends at ‘He’ because ‘He’ has complete duplet.
distribution of electrons in an atom?
Electrons are filled in the second shell
When Mendeleev put forth his along the second period. The second period,
periodic table in 1869, the atomic structure thus, begins with Li (Z=3) : 1s2 2s1 and ends
was not known. He observed periodicity in up with Ne (Z=10): 1s22s22p6. ‘Ne’ with 8
the properties of elements on arranging them electrons in its outermost second shell has
in an increasing order of atomic mass. Later, complete octet. The second shell has electron
with the advent of quantum mechanical model capacity of 8. It gets filled along the second
of atom, the properties of elements were period, as the atomic number increases. Thus
correlated to electronic configuration. there are eight elements in the second period.
You have learnt in the Chapter 4
that the electrons in atom are distributed in Similarly there are eight elements Na
shells and subshells in accordance with the (Z=11) to Ar (Z=18) having the condensed
aufbau principle which includes increasing electronic configurations described together as
order of energy, Pauli exclusion principle [Ne]3s1-23p1-6 in the third period, as a result of
and Hund’s rule of maximum multiplicity. completion of the third shell.
When elements are arranged in an increasing
order of atomic number (Z), periodicity is The fourth period begins with filling
observed in their electronic configuration and of 4s subshell. The first two elements of the
which reflects in the characteristic structure fourth period are K (Z=19) : [Ar] 4s1 and Ca
of the modern periodic table. The location (Z=20): [Ar]4s2. According to the aufbau
of elements in the modern periodic table is principle the next higher energy subshell is 3d,
correlated to quantum numbers of the last which can accommodate upto 10 electrons.
filled orbital. Let us have a deeper look into Filling of the 3d-subshell results in the next 10
the electronic configuration of the elements elements of the fourth period, from Sc (Z=21) :
and the structure of the modern periodic table. [Ar] 4s23d1 to Zn (Z=30) : [Ar] 4s23d10. After
7.3.1 Electronic configuration in periods this the electrons enter the subshell 4p for
the next six elements : Ga (Z = 31) : [Ar]
We noted earlier that periods in the 4s23d104p1 to Kr (Z=36) : [Ar] 4s23d104p6. The
modern periodic table are numbered 1 to 7. fourth period, thus, contains in all 18 elements
On inspection of the electronic configurations (2+10+6=18).
(see Fig. 7.3) of elements in various periods The fifth period accomodates 18
we understand that the period number is same elements as a result of successive filling of
as the principal quantum number ‘n’ of the electrons in the 5s, 4d and 5p subshells.
outermost or valence shell of the elements.

95

Representative elements Representative elements
s - Block p - Block
1

d - Block

96

f - Block Lanthanoid
4f n 5d0-1 6s2

Actinoids
5f n 6d0-1 7s2

Fig. 7.2 : Outer electronic configuration of elements in the four blocks of the periodic table

Problem 7.1 : What is the subshell In short, a period begins by filling of one
in which the last electron of the first electron to the ‘s’ subshell of a new shell and
ends with an element having complete octet
element in the 6th period enters ? (or duplet) corresponding to the same shell.
Solution : Between these two ends corresponding to ‘s’
The 6th period begins by filling the and ‘p’ subshell of the valence shell, the inner
subshells ‘d’ and ‘f’ are filled successively
last electron in the shell with n=6. The
following the aufbau principle.
lowest energy subshell of any shell is
‘s’. Therefore the last electron of the 7.3.2 Electronic configuration in groups
first element in the 6th period enters the
subshell ‘6s’. A new shell is added down a group.
The general outer electronic configuration,
Problem 7.2 : How many elements are therefore, is expected to be the same down any
present in the 6th period? Explain. particular group. Indeed it is found to be so for
the groups 1, 2 and 3. (See Fig. 7.2 and Table
Solution : 7.1) In the groups 13 to 18 the appropriate
The 6th period begins by filling the last inner ‘d’ and ‘f’ subshells are completely filled
electron in the subshell ‘6s’ and ends by and the general outer electronic configuration
completing the subshell ‘6p’. Therefore, is the same down the groups 13 to 18. (see Fig.
the sixth period has the subshells filled
in increasing order of energy as 6s < 4f < 7.2 and Table 7.1).
5d < 6p. The electron capacities of these In the groups 4 to 12, however, the ‘d’
subshells are 2, 14, 10 and 6, respectively.
Therefore, the total number of elements in and ‘f’ subshells are introduced at a later stage
the 6th period are 2+14+10+6 = 32. (4th period for ‘d’ and 6th period for ‘f’) down
the group. As a result variation in the general
outer configuration is introduced only at the

later stage down the groups 4 to 12.

Table 7.1 : General outer electronic configuration in groups 1 to 3 and 13 to 18

Group number General outer configuration Examples

Group 1 ns1 3Li:2s1, 11Na:3s1
Group 2 ns2 4Be:2s2, 12Mg:3s2
Group 3 ns2 (n-1)d1 21Sc:4s23d1, 39Y:5s24d1
Group 13 ns2 np1 5B:2s22p1, 13A1:3s23p1
Group 14 ns2 np2 6C:2s22p2, 14Si:3s23p2
Group 15 ns2 np3 7N:2s22p3, 15P:3s23p3
Group 16 ns2 np4 8O:2s22p4, 16S:3s23p4
Group 17 ns2 np5 9F:2s22p5, 17C1:3s23p5
Group 18 ns2 np6 10Ne:2s22p6, 18Ar:3s23p6

97

7.3.3 Electronic configuration in the four ten groups, namely, groups 3, 4, 5, 6, 7, 8, 9,
blocks: We noted in section 7.2 that structure 10, 11 and 12 in the d-block which appears
of the modern periodic table shows four blocks. in the centre of the modern periodic table.
These blocks are formed in accordance with The general outer electronic configuration of
the subshell in which the last electron enters. the d-block elements is ns0-2 (n-1)d1-10. Some
Accordingly the four blocks are named as variations in the configuration, consequent
s-block, p-block, d-block and f-block. to the extra stability associated with half-
filled and a fully filled subshell, are readilly
s-Block : The last electron in the s-block observed. For example, the outer electronic
elements is filled in a s-subshell. There being configuration of Cr (Z = 24) is 4s13d5 instead
only one orbital in a s-subshell, the general of 4s23d4. This is because both 4s and 3d
outer electronic configuration of s-block subshells are half-filled.
f-Block : In the f-block elements the last
elements is ns1-2. Thus elements of the group-1 electron is filled in f-orbital. As there are
and group-2 belong to the s-block. The s-block seven orbitals in a f-subshell, the general outer
is present on the left extreme of the modern electronic configuration of the f-block is ns2
(n-1)d0-1 (n-2)f1-14. The variations as a result
periodic table. of the extra stability of half-filled and fully
filled subshell need to be accounted for. For
p-Block: The last electron in the p-block example, the outer electronic configuration
elements is filled in p-subshell. There being of 63Eu is 6s24f75d0 because the 4f subshell is
three degenerate p-orbitals in a p-subshell, half-filled. The f-block constitutes two series
upto 6 electrons can be filled. Therefore, the of 14 elements called the lanthanoid and the
elements belonging to six groups, namely, actinoid series, put one below the other. The
group 13, 14, 15, 16, 17 and 18 constitute the f-block is placed separately at the bottom of
p-block. The p-block appears on the right in the periodic table.
the modern periodic table. The p-block ends
with the group 18 which represent the family Problem 7.3 : Outer electronic configu-
of inert gases. Remarkably, the first element of rations of a few elements are given be-
the group 18, helium (He) does not have the p low. Explain them and identify the period,
group and block in the periodic table to
subshell as its valence shell has n =1 and its which they belong.
2He : 1s2, 54Xe : 5s25p6, 16S : 3s23p4, 79Au :
configuration is shown as 1s2. Yet ‘He’ is placed 6s15d10
in the 18th group of the p –block because its Solution :
valence shell is completely filled (which is a 2He : 1s2
complete duplet), similar to complete valence Here n = 1. Therefore, 2He belongs to the
shell of the other elements belonging to group 1st period. The shell n =1 has only one
18 (which have complete octets). The general subshell, namely 1s. The outer electronic
electronic configuration for the p-block (from configuration 1s2 of ‘He’ corresponds to
the maximum capacity of 1s, the complete
the second to the seventh period) is ns2 np1-6. duplet. Therefore, He is placed at the end
of the 1st period in the group 18 of inert
d-Block : The d-block in the modern periodic
table is formed as a result of filling the last
electron in d-orbital. A d-subshell is present in

the shells with n ≥ 3 and according to the (n+l)
rule (refer to Chapter 4) the energy of ns orbital

is less than that of the (n-1)d orbital. As a result,

the last electron enters a (n-1)d-subshell only

after the ns subshell is completely filled. There

being five orbitals in a d-subshell, 10 electrons
can successively be accommodated. There are

98

gases, so ‘He’ belongs to p-block. 7.4.2 Characteristics of p-block elements :
54Xe : 5s25p6 The p-block contains elements of groups

Here n = 5. Therefore, 54Xe belongs to the 5th 13 to 18. The p-block elements together
period. The outer electronic configuration
5s25p6 corresponds to complete octet. with s-block elements are called main group
Therefore 54Xe is placed in group 18 and
belongs to p-block. elements or representative elements. The last
16S : 3s23p4
Here n = 3. Therefore, 16S belongs to the 3rd group of the p-block, namely, the 18th group,
period. The 3p subshell in ‘S’ is partially
filled and short of completion of octet by is the family of noble/inert gases. These have
two electrons. Therefore ‘S’ belongs to (18- closed valence shells (complete duplet in the
2) = 16th group and p – block. case of ‘He’ and complete octet in the case of
79Au : 6s15d10 the other noble gases) and therefore very low
Here n = 6. Therefore, ‘Au’ belongs to the
6th period. The sixth period begins with chemical reactivity. Elements of group 17
filling of electron into 6s and then into 5d (halogen family) and group 16 (chalcogens)
orbital. The outer configuration of ‘Au’ :
6s1 5d10 implies that (1+10) = 11 electrons include reactive nonmetals. The electron
are filled in the outer orbitals to give ‘Au’.
Therefore ‘Au’ belongs to the group 11. As gain enthalpies being highly negative, they
the last electron has entered ‘d’ orbital ‘Au’ gain one or two electrons readily and form
belongs to the d-block. anions (X or X2 ) which have complete

7.4 Blockwise characteristics of elements octet. The p-block contains all the three
We have seen in the previous section
traditional types of elements. The metals
that the 118 elements in the modern periodic on the left, the nonmetals on the right and the
table are distributed in four blocks having metalloids along a zig-zag line (see Fig. 7.1)
general electronic configuration according which separates metals from nonmetals. The
to their block. The elemental properties are nonmetallic character increases as we move
characteristic of the block they belong. from the left to the right, whereas it decreases
7.4.1 Characteristics of s-block elements as we go down a group.

The s-block contains the elements of 7.4.3 Characteristics of d-block elements
group 1 (alkali metals) and group 2 (alkaline
earth metals). All these elements are reactive The d-block contains elements of
metals, and occur in nature only in combined the groups 3 to 12. They are all metals. The
state. Their compounds, with exception of d-block elements are known as transition
Li and Be, are predominantly ionic. This is elements or transition metals. They form a
because they have only one or two valence bridge between chemically reactive s-block
electrons which they can lose readily forming elements and less reactive elements of groups
M⊕ or M2⊕ ions. They have low ionization 13 and 14. Most of d-block elements possess
enthalpies, which decrease down the group Partially filled inner d-orbitals. As a result
resulting in increased reactivity. the d-block elements have properties such as
variable oxidation state, paramagnetism,
ability to form coloured ions. They can
be used as catalysts. Zn, Cd, and Hg with
configuration ns2 (n-1) d10, (completely filled s
- and d - subshells) do not show the properties
typical of transition metals.
7.4.4 Characteristics of f-block elements

The f-block contains elements all of which

are metals and are placed in the two rows called

lanthanoid series (58Ce to 71Lu) and actinoid

99

series (90Th to 103Lr). These series are named properties. We will now discuss in this section
after their preceding elements lanthanum the periodic trends in some physical and
(57La) and actinium (89Ac) in the third group of chemical properties of elements, correlating
the d-block of the 6th and 7th period respectively. them with electronic configuration and the
The lanthanides are also known as rare earth nuclear charge. These trends are explained
elements. The last electron of the elements of in terms of two fundamental factors, namely,
these series is filled in the (n-2)f subshell, and attraction of extranuclear electrons towards
therefore, these are called inner-transition the nucleus and repulsion between electrons
elements. These elements have very similar belonging to the same atom. These attractive
properties within each series. The actinide and repulsive forces operate simultaneously
in an atom. This results in two interrelated
Problem 7.4 : Chlorides of two metals are phenomena called effective nuclear charge
common laboratory chemicals and both and screening effect.
are colourless. One of the metals reacts
vigorously with water while the other re- 7.5.1 Effective nuclear charge and screening
acts slowly. Place the two metals in the effect : In a multi-electron atom the positively
appropriate block in the periodic table. charged nucleus attracts the negatively
justify your answer. charged electrons around it, and there is mutual
Solution repulsion amongst the negatively charged
Metals are present in all the four blocks of extranuclear electrons. The repulsion by inner-
the periodic table. Salts of Metals in the shell electrons is particularly important. This
f-block and p-block (except AlCl3) are not results in pushing the outer-shell electrons
common laboratory chemicals. Therefore, further away from the nucleus. The outer-
the choice is between s- and d-block. From shell electrons are, thus, held less tightly by
the given properties their placement is the nucleus. In other words, the attraction of
done as shown below: the nucleus for an outer electron is partially
s-block: Metal that reacts vigorously with cancelled. It means that an outer-shell electron
water does not experience the actual positive charge
d-block: Metal that reacts slowly with present on the nucleus. The net nuclear charge
water. actually experience by an electron is called the
The colourless nature of the less effective nuclear charge, Zeff. The effective
reactive metal in the d-block implies that nuclear charge is lower than the actual nuclear
the inner d-subshell is completely filled. charge, Z. In other words, the inner electrons
shield the outer electrons from the nucleus to a
elements beyond 92U are called transuranium certain extent. This effect of the inner electrons
elements. All the transuranium elements are on the outer electrons is called screening effect
manmade and radioactive. or shielding effect of the inner/core electrons.
7.5 Periodic trends in elemental properties
Effective nuclear charge = Zeff
The original structure of the periodic table = Z - electron
was based on empirically observed periodicity
in the elemental properties. Thus, elemental shielding
properties show similarity in a group and =Z- s
show gradual variation across a period. The
quantum mechanical model of atom explains Here s (sigma) is called shielding
the observed periodic trends of elemental constant or screening constant and the value
of s depends upon type of the orbital that the
electron occupies.

100

As we move across a period actual is estimated to be 77 pm. (ii) Bond length
nuclear charge increases by +1 at a time, the of Cl-Cl bond in Cl2 is measured as 198 pm.
valence shell remains the same and the newly Therefore, the atomic radius of Cl is estimated
added electron gets accommodated in the to be 99 pm. (see Fig. 6.3)
same shell. There is no addition of electrons
to the core. Thus, shielding due to core 224 pm
electrons remains the same though the actual
nuclear charge, Z, increases. The net result is Be Be Be
that the effective nuclear charge, Zeff, goes
on increasing across a period. On the other Be Be Be
hand, the Zeff decreases down a group. This
is because, as we move down a group, a new Be Be Be
larger valence shell is added. As a result, there
is an additional shell in the core. The shielding Atomic radius of Be = 224 pm = 112 pm
effect of the increased number of core electrons 2
outweighs the effect of the increased nuclear
charge; and thereby the effective nuclear 198 pm Cl
charge felt by the outer electrons decreases
largely down a group. Cl
7.5.2 Periodic trends in physical properties
Cl2 = 198 pm = 99 pm
Many physical properties of elements Atomic radius of Cl 2
such as melting point, boiling point and density
show periodicity. In this section we are going Fig 6.3 : Estimation of atomic radii of metals
to consider the periodic trends in physical
properties with reference to atomic radius, and nonmetals
ionic radius, ionization enthalpy, electron gain
enthalpy and electronegativity. In the case of metals, distance between
a. Atomic radius : You have learnt in chapter the adjacent atoms in metallic sample is
4 that the quantum mechanical model of atom measured. One half of this distance is taken
describes the extranuclear part of atom as as the metallic radius. Thus, atomic radius is
the electron cloud. As a direct implication of one half of the internuclear distance between
this, the atom has no definite boundary. The two adjacent atoms of a metal or two single
atomic size or atomic radius, therefore, can bonded atoms of a nonmetal.
be estimated from the internuclear distance
under different circumstances. In the case of Do you know ?
nonmetals (except noble gases), the atoms
of an element are bonded to each other by Atomic radius is estimated in terms of
covalent bonds. (Refer to Chapter 5). Bond the electron density surface which encloses
length of a single bond is taken as sum of radii typically 95 % (or more, which is orbitary
of the two single bonded atoms. This is called of the electron density.
covalent radius of the atom. For example
: (i) Bond length of C-C bond in diamond is
154 pm. Therefore, atomic radius of carbon

101

Table 7.2 : Atomic radii of some elements
Element symbol (atomic radius / pm)

Group 1 2 13 14 15 16 17
Period
Be (111) B (88) C (77)
2 Li (152) Mg (160) Al (143) Si (117) N (74) O (66) F (64)
P (110) S (104) Cl (99)
3 Na (186) Br (114)
I (133)
4 K (231) At (140)

5 Rb (244)

6 Cs (262)

It can be seen from Table 7.2 that atomic An anion has a larger radius than the
radius decreases across a period (upto group corresponding atom, as it has more number
17) and increases down a group. Greater the of electrons than the atom. These additional
effective nuclear charge stronger is attraction electrons result in increased electron repulsion,
of the nucleus for the outer electrons and decreased effective nuclear charge and in turn,
smaller is the atomic radius. As we move across the increased size.
a period, screening effect caused by the core
electrons remains the same, on the other hand Some atoms and ions contain the same
the effective nuclear charge goes on increasing number of electrons and are called isoelectronic
(see section 6.5.1). The valence electrons are, species. (See chapter 4) The actual nuclear
therefore, more tightly bound and in turn the charge of the isoelectronic species is, however,
atomic radius goes on decreasing along a different. The radii of isoelectronic species
period. The effective nuclear charge decreases vary according to actual nuclear charge. Larger
and shielding effect increases down a group nuclear charge exerts greater attraction for the
(see section 7.5.1). The valence electrons are, electrons and the radius of that isoelectronic
thus, held by weaker attractive force and the species becomes smaller. For example, F and
atomic radius increases down a group. Na⊕ both have 10 electrons, their radii are 133
b. Ionic radius: An atom forms a positively pm and 98 pm respectively, as the nuclear
charged ion, cation, on the removal of one or charge of F is + 9 which is smaller than that of
more electrons whereas a negatively charged Na⊕ which is, + 11.
ion, anion, is formed with gain of one or c. Ionization enthalpy : Removal of an
more electrons. Measurements of distances electron from the neutral atom X results
between neighbouring cations and anions in
ionic crystals have been useful for estimation Problem 7.5 : Identify the species having
of ionic radii. The ionic radii show the same larger radius from the following pairs : (i)
trends as of atomic radii. Na and Na⊕, (ii) Na⊕ and Mg2⊕
Solution :
A cation is smaller than the atom from (i) The nuclear charge is the same in Na
which it is formed, because it contains fewer and Na⊕. But Na⊕ has less number of elec-
electrons than atom, though the nuclear charge trons and less number of occupied shells
is the same. As a result the shielding effect (two shells in Na⊕ while three shells in Na).
is less and effective nuclear charge is larger Therefore, radius of Na is larger. (ii) Na⊕
within a cation. and Mg2⊕ are isoelectronic species. Mg2⊕
has larger nuclear charge than Na⊕. There-
fore, Na⊕ has larger radius.

102

in formation of cation X⊕. The energy Table 7.3 : First ionization enthalpy values of

required to remove an electron from elements of group 1

the isolated gaseous atom in its ground Elements Atomic Outer ∆iH1
of Group Number electronic /kJ
state is called ionization enthalpy (∆iH). configuration mol-1
Ionization enthalpy is the quantitative 1Z

measure of tendency of an element to lose Li 3 2s1 520

electron and expressed in kJ mol-1. Since Na 11 3s1 496

electrons are lost one at a time, we have K 19 4s1 419

first ionization enthalpy, second ionization Rb 37 5s1 403

enthalpy, and so on, for a given element. Cs 55 6s1 374

X (g) X⊕ (g) + e ; ∆iH1 An overall increase of the first ionization

X⊕ (g) X2⊕(g) + e ; ∆iH2 enthalpy across the period 2 (see table 7.4)

Ionization enthalpy is always positive can be notice. This is because the screening

since the energy always need to be supplied to is the same and the effective nuclear charge

knock out electron from atom. increases across a period. (See section 7.5.1)

The second ionization enthalpy, ∆iH2, is As a result the outer electron is held more
larger than the first ionization enthalpy, ∆iH1,
as it involves removal of electron from the tightly. The first ionization enthalpy, therefore,
positively charged species. Tables 7.3 and 7.4
show the values of first ionization enthalpy increases across a period. The alkali metal
down the group 1 and across the period 2
respectively. displays the lowest first ionization enthalpy.

It is seen from Table 7.3 that on The inert gas shows the highest first ionization
moving down the group the first ionization
enthalpy decreases. This is because electron enthalpy across a period.
is to be removed from the larger valence
shell. Screening due to core electrons goes on Some irregularities are noticed for the
increasing and the effective nuclear charge
decreases down the group. (See section 7.5.1) first ionization enthalpies as we move across
The removal of the outer electron, therefore,
becomes easier. a period. For example, the first ionization

enthalpy of ‘B’ is smaller than that of ‘Be’.

This is because ‘Be’ loses the electron from

2s orbital while ‘B’ loses the electron from

2p orbital which has less penetration (see

shapes of orbitals in chapter 4) than the 2s

orbital and therefore it is easier to remove a

2p electron than a 2s electron. Similarly, first

ionization enthalpy of ‘O’ is smaller than that

of ‘N’. This is because ‘O’ loses the electron

from a doubly occupied ‘2p’ orbital. Due to

Table 7.4 : First ionization enthalpy values of elements of period 2

Element of Li Be B C N O F Ne

period 2

Atomic 3 4 5 6 7 8 9 10

number

Z

Outer 2s1 2s2 2s2 2p1 2s2 2p2 2s2 2p3 2s2 2p4 2s2 2p5 2s2 2p6

electronic

Configuration

∆iH1 / 520 899 801 1086 1402 1314 1681 2080
kJ mol-1

103

electron-electron repulsion it is easier to lose the other the hand, exhibit high positive values
this electron than an electron from the singly for electron gain enthalpy. Since, the added
occupied 2p orbital in nitrogen atom. (This is electron has to enters the next higher shell
a consequence of Hund’s rule of maximum with larger principal quantum number, this is
multiplicity, described in Chapter 4.) very unstable electronic configuration due to
very low effective nuclear charge and high
Problem 7.6 : shielding from the core electrons. The alkali
The first ionization enthalpy values of metals have very low negative electron gain
Si, P and C1 are 780, 1060 and 1255 kJ enthalpy values. In general, electron gain
mo1-1 respectively. Predict whether the first enthalpies are large negative for elements of
ionization enthalpy of S will be closer to the upper right of the periodic table, excluding
1000 or 1200 kJ mol-1. the group 18 of noble gases.
Solution : The trends for electron gain enthalpy values
As we move across the period 3 from left of the elements in the periodic table are less
to right the elements Si, P, S, C1 come regular than the ionization enthalpies. An
in a sequence. Their outer electronic overall trend reveals that electron gain
configurations are 3s23p2, 3s23p3, 3s23p4 and enthalpy becomes more negative with increase
3s23p5 respectively and ‘P’ loses an electron of atomic number along the period upto the
from a singly occupied 3p orbital whereas second last element, because the effective
‘S’ loses an electron from a doubly occupied nuclear charge increases across the period and
3p orbital. Therefore, the first ionization it is easier to add an electron to a smaller atom.
enthalpy value of ‘S’ has to be lower than Down the group, the electron has to be add to
that of ‘P’. Thus, first ionization enthalpy farther shell and the electron gain enthalpy,
of ‘S’ would be less than 1060 kJ mol-1, and thus, becomes less negative.
therefore, should be close to 1000 kJ mol-1
and not 1200 kJ mo1-1. Problem 7.7 :

d. Electron gain enthalpy: Addition of Identify the element with more negative
an electron to a neutral atom (X) results in
formation of an anion (X ). The enthalpy value of electron gain enthalpy from the
change that takes place when an electron
is added to an isolated gaseous atom in its following pairs. Justify.
ground state is called the electron gain
enthalpy, ∆egH. Electron gain enthalpy is a (i) C1 and Br (ii) F and O
quantitative measure of the ease with which an
atom adds an electron forming the anion and is Solution :
expressed in units of kJ mol-1.
(i) CI and Br belong to the same group of
X(g) + e X (g); ∆egH.
∆egH may be positive or negative halogens with Br having higher atomic
depending upon whether the process of
adding electron is endothermic or exothermic. number than Cl. As the atomic number
Elements of group 17 have very high negative
values of electron gain enthalpy. This is because increases down the group the effective
of they attain stable noble gas electronic
configuration on the addition of one electron. nuclear charge decreases. The increased
Elements of group 18, noble gases have, on
shielding effect of core electrons can be

noticed. The electron has to be added to

a farther shell, which releases less energy

and thus electron gain enthalpy becomes

less negative down the group. Therefore, Cl

has more negative electron gain enthalpy

than Br. ....CONTD on next page

104

(ii) F and O belong to the same second electron involved in formation of the covalent
period, with F having higher atomic bond. Electronegativity increases as we move
number than O. As the atomic number across the period. This is because the effective
increases across a period, atomic radius nuclear charge increases steadily across the
decreases, effective nuclear charge period. The EN decreases down the group. The
increases and electron can be added more size of the valence shell goes on increasing,
easily. Therefore more energy is released the shielding effect of the core electron goes
with gain of an electron as we move to on increasing and in term the effective nuclear
right in a period. Therefore, F has more charge decreases down the group.
negative electron gain enthalpy.
Electronegativity predicts the nature of the
e. Electronegativity: When two atoms of bond, or, how strong is the force of attraction
different elements form a covalent bond, that holds two atoms together. (Refer Chapter
the electron pair is shared unequally. The 5).
ability of a covalently bonded atom to
attract the shared electrons toward itself 7.5.3 Periodic trends in chemical properties:
is called electronegatively (EN). It is not an The most fundamental chemical property
experimentally measurable quantity. A number
of numerical scales of electronegativity were of an element is its combining power. This
developed by many scientists. Pauling scale of property is numerically expressed in terms of
electronegativity is the one used most widely. valency or valence. Valency of an element
Linus Pauling assigned (1922) arbitrarily a indicates the number of chemical bonds that
value of 4.0 for fluorine which is expected to the atom can form giving a molecule. Another
have the highest value of electronegativity.
Table 7.5 shows the values of electronegativity frequently used term related to valence is the
assigned to some other elements. oxidation state or oxidation number. Valence
does not have any sign associated with it, but
Electronegativity represents attractive oxidation number does, and can be either + or
force exerted by the nucleus on shared –, which is decided by electronegativities of
electrons. Electron sharing between covalently atoms that are bonded.
bonded atoms takes place using the valence
electron. The electronegativity depends upon The second aspect of chemical property
the effective nuclear charge experienced by is the chemical reactivity. The chemical
reactivity is related to the ease with which an
element loses or gains the electrons.

The chemical properties of elements are

relate to electronic configuration.

Table 7.5 : Electronegativity (EN) values of some elements (Pauling scale)
Element symbol (EN)

Group 1 2 13 14 15 16 17
Period

1 H (2.1)

2 Li (1.0) Be (1.5) B (2.0) C (2.5) N (3.0) O (3.5) F (4.0)

3 Na (0.9) Mg (1.2) Al(1.5) Si (1.8) P (2.1) S (2.5) Cl (3.0)

4 K (0.8) Br (2.8)

5 Rb (0.8) I (2.5)

6 Cs (0.7) At (2.2)

105

Table 7.6 : Periodic trends in valency of main group elements

Group 1 2 13 14 15 16 17 18
ns2np5 ns2np6
General outer electronic ns1 ns2 ns2np1 ns2np2 ns2np3 ns2np4
configuration 78

Number of valence 123 4 5 6 1,7 0
electrons HF
1 2 3 4 3,5 2,6 HCI
Valency HBr
OF2
Formula of hydride LiH BeH2 B2H6 CH4 NH3 H2O
CI2O7
NaH MgH2 AlH3 SiH4 PH3 H2S
Br2O
Formula of oxide KH CaH2 GaH3 GeH4 AsH3 H2Se
Li2O BeO B2O3 CO2
N2O3 SO2, SO3
Na2O MgO Al2O3 SiO2 N2O5 SeO2, SO3
P4O6,
K2O CaO Ga2O2 GeO2 P2O5,
As2O3,
As2O5

a. Periodic trends in valency: Valency of the The ionization enthalpy decreases down the

main group elements is usually equal to the group. The tendency to lose valence electrons,
number of valence electrons (outer electrons) thus, increases down the group and the
and/or equal to difference between 8 and the metallic character increases down a group.
number of valence electrons. Table 7.6 shows The ionization enthalpy increases across the
the periodic trends in the valency of main group period and consequent metallic character
elements by taking examples of hydrides and decreases across a period. Electron gain
some oxides. enthalpy becomes more and more negative
across the period, it tends to be less negative
As may be notice form Table 7.6 that down the group. Thus, nonmetallic character
the valency remains the same down the group increases across the period and decreases
and shows a gradual variation across the
period as atomic number increases from left to down the group.

right. c. Periodic trends in chemical reactivity:

b. Periodic trends in metallic–nonmetallic Chemical reactivity of elements is decided by
character : The metals, nonmetals and how easily it attains electronic configuration
metalloids appear in separate regions of modern of the nearest inert gas by gaining or loosing
periodic table: metals on the left, nonmetals electrons.
on the right and metalloids along a zig-zag line
separating the two. Metals are characterized The elements preceding an inert gas
by good electrical conductivity and ability to react by gaining electrons in the outermost
form compounds by loss of valence electrons. shell, whereas the elements which follow an
Nonmetals are characterized by their poor inert gas in the periodic table react by loss of
electrical conductivities and ability to form valence electrons. Thus the chemical reactivity
compounds by gain of valence electrons in is decided by the electron gain enthalpy and
valence shell. This can be explained in terms of ionization enthalpy values, which in turn, are
ionization enthalpy and electron gain enthalpy. decided by effective nuclear charge and finally
by the atomic size. The ionization enthalpy is
smallest for the element on the extreme left in

106

a period, whereas the electron gain enthalpy is Problem 7.8 :
most negative for the second last element on Ge, S and Br belong to the groups 14, 16
the extreme right, (preceding to the inert gas and 17, respectively. Predict the empirical
which is the last element of a period). Thus, formulae of the compounds those can be
the most reactive elements lie on the extreme formed by (i) Ge and S, (i) Ge and Br.
left and the extreme right (excluding inert Solution :
gases) of the periodic table. From the group number we understand that
the general outer electronic configuration
Apart from the metallic-nonmetallic and number of valence electrons and
character the chemical reactivity can be valencies of the three elements are :
illustrated by comparing their reaction with
oxygen to form oxides and the nature of the Element Group Outer Number Valency
oxides. The reactive elements on the extreme electronic of
left (that is, alkali metals) react vigorously configuration Valence
with oxygen to form oxides (for example electron
Na2O) which reacts with water to form strong
bases (like NaOH). The reactive elements on Ge 14 ns2np2 4 4
the right (that is, halogens) react with oxygen
to form oxides (for examples Cl2O7) which on S 16 ns2np4 6 8-6
reaction with water form strong acids (like
HCIO4). The oxides of the elements in the =2
center of the main group elements are either
amphoteric (for example Al2O3), neutral Br 17 ns2np5 7 8-7
(for example CO, NO) or weakly acidic (for
example CO2) =1

The change in atomic radii of transition (i) S is more electronegative than Ge.
metals and inner transition elements is rather
small. Therefore, the transition metals and Therefore, the empirical formula of the
inner transition elements belonging to the
individual series have similar chemical compound formed by these two elements
properties. Their ionization enthalpies are
intermediate between those of s-block and is predicted by the method of cross
p - block.
multiplication of the valencies :

Element: Ge S

Valency: 4 2

Formula: Ge2S4
Empirical formula: GeS2
(ii) Br is more electronegative than Ge. The

empirical formula of the compound formed

by these two elements in predicted by the

method of cross multiplication of valencies :

Element: Ge Br

Valency: 4 1

Formula: GeBr4
Empirical formula: GeBr4

107

Problem 7.9
Write the chemical equations for reaction, if any, of (i) Na2O and (ii) Al2O3 with HCl and NaOH
both. Correlate this with the position of Na and Al in the periodic table, and infer whether the
oxides are basic, acidic or amphoteric.
Solution
(i) Na2O + 2HCl 2NaC1 + H2O
Na2O + NaOH No reaction
As Na2O reacts with an acid to form salt and water it is a basic oxide. This is because
Na is a reactive metal lying on the extreme left of the periodic table.
(ii) Al2O3 + 6HCI 2AlCl3 + 3H2O
Al2O3⊕ 2NaOH 2Na AlO2 + H2 O
As Al2O3 reacts with an acid as well as base to form a salt and water. It is an amphoter-
ic oxide. Al is a moderately reactive element lying in the centre of main group elements in the
periodic table.

Exercises

1. Explain the following D. Why the second ionization enthalpy is
greater than the first ionization enthalpy
A. The elements Li, B, C, Be and N have ?

the electronegativities 1.0, 2.0, 1.5 and E. Why the elements belonging to the same
group do have similar chemical
3.0, respectively on the Pauling scale. properties ?

B. The atomic radii of Cl, I and Br are 99, F. Explain : electronegativity and electron
gain enthalpy. Which of the two can be
133 and 114 pm, respectively. measured experimentally?

C. The ionic radii of F and Na⊕ are 133 4. Choose the correct option
A. Consider the elements B, Al, Mg and K
and 98 pm, respectively. predict the correct order of metallic
character :
D. C1135APu lifsoisramansomncomeltoeatula,rle1.4dSsiailstsawmhielteaZllonidforamnds a. B > Al > Mg > K
E. b. Al > Mg > B > K
c. Mg > Al > K > B
colourless salts. d. K > Mg > Al > B
B. In modern periodic table, the period
2. Write the outer electronic configuration number indicates the :
a. atomic number
of the following using orbital notation b. atomic mass
c. principal quantum number
method. Justify. d. azimuthal quantum number

A. Ge (belongs to period 4 and group 14)

B. Po (belongs to period 6 and group 16)

C. Cu (belongs to period 4 and group 11)

3. Answer the following

A. La belongs to group 3 while Hg belongs

to group 12 and both belong to period 6

of the periodic table. Write down the

general outer electronic configuration of

the ten elements from La to Hg together

using orbital notation method.

B. Ionization enthalpy of Li is 520 kJ mol-1

while that of F is 1681 kJ mol-1. Explain.

C. Explain the screening effect with a

suitable example.

108

C. The lanthanides are placed in the D. With the help of diagram answer the
questions given below:
periodic table at

a. left hand side

b. right hand side Li Be O
S
c. middle

d. bottom

D. If the valence shell electronic

configuration is ns2np5, the element will

belong to

a. alkali metals

b. halogens a. Which atom should have smaller

c. alkaline earth metals ionization enthalpy, oxygen or sulfur?

d. actinides b. The lithium forms +1 ions while

E. In which group of elements of the modern berylium forms +2 ions ?

periodic table are halogen placed ? E. Define : a. Ionic radius

a. 17 b. 6 b. Electronegativity

c. 4 d. 2 F. Compare chemical properties of metals

F. Which of the atomic number represent and non metals.

the s-block elements ? G. What are the valence electrons ? For

a. 7, 15 s-block and p-block elements show that

b. 3, 12 number of valence electrons is equal to

c. 6, 14 its group number.

d. 9, 17 H. Define ionization enthalpy. Name the

G. Which of the following pairs is NOT factors on which ionisation enthalpy

isoelectronic ? depends? How does it vary down the

a. Na⊕ and Na group and across a period?

b. Mg2⊕ and Ne I. How the atomic size vary in a group and

c. Al3⊕ and B3⊕ across a period? Explain with suitable

d. P3 and N3 example.

H. Which of the following pair of elements J. Give reasons.

has similar properties ? a. Alkali metals have low ionization

a. 13, 31 enthalpies.

b. 11, 20 b. Inert gases have exceptionally high

c. 12, 10 ionization enthalpies.

d. 21, 33 c. Fluorine has less electron affinity than

5. Answer the following questions chlorine.

A. The electronic configuration of some d. Noble gases possess relatively large

elements are given below: atomic size.

a. 1s2 b. 1s2 2s2 2p6 K. Consider tohxeidoexiwdeosuLldi2Oyo, uCOex2p, eBc2tOt3o.
a. Which
In which group and period of the be

periodic table they are placed ? the most basic?

B. For each of the following pairs, indicate b. Which oxide would be the most acidic?

which of the two species is of large size : c. Give the formula of an amphoteric

a. Fe2+ or Fe3+ b. Mg2+ or Ca2+ oxide.

C. Select the smaller ion form each of the

following pairs:

a. K+ , Li+ b. N3-, F- Activity :

Prepare a wall mounting chart of the
modern periodic table.

109

8. Elements of Groups 1 and Group 2

We have seen in chapter 7 that the element Electronic configuration of hydrogen is
of group 1 and group 2 belong to the s-block of
the modern periodic table. 1s1 which is similar to ns1 which is the outer

Can you recall? electronic configuration of alkali metals

1. Which is the first element in the belonging to the group 1. But 1s1 also resembles
periodic table ?
the outer electronic configuration of group
2. What are isotopes ?
3. Write the formulae of the compounds 17 elements, which is ns2np5. By adding one

of hydrogen formed with sodium and electron to H we get electronic configuration
chlorine.
of the inert gas He which is 1s2 and by
Hydrogen is the first element in the
periodic table. Hydrogen appears at the top of adding one electron to ns2np5 we get ns2np6
group 1 of the alkali metals. But it differs from
alkali metals in a number of respects and, which is the outer electronic configuration
therefore, is studied separately.
8.1 Hydrogen : Hydrogen has the simplest of the remaining inert gases. Some chemical
atomic structure of all the elements. Ahydrogen
atom consists of a nucleus of charge +1 and properties of hydrogen are similar to those of
one extranuclear electron. Hydrogen has a
little tendency to lose this electron however alkali metals while some resemble halogens.
can pair with the other electron easily forming
a covalent bond. It exists in diatomic form The uniqueness of hydrogen is that H⊕ formed
as H2 molecule; therefore it is often called
dihydrogen. by loss of the electron from hydrogen atom
8.1.1 Occurrence : In the free state hydrogen
exists as dihydrogen gas. Hydrogen is most does not exist freely. It is always associated
abundant element in the universe; it makes 70
% of the total mass of the universe. Hydrogen with other molecules. For example:
is also the principal element in the solar
system. On the earth, hydrogen is the tenth H⊕ + H2O H3O⊕
most abundant element on mass basis and the This is because H⊕ is nothing else but
third most abundant element on atom basis.
8.1.2 Position of hydrogen in the periodic a proton. Hydrogen is, therefore, placed
table : Position of hydrogen in the periodic
table has been a subject of discussion. separately above the group 1. It may be noted,

here, that metastable metallic hydrogen was

discovered at Harvard university, USA, in

January 2017.

8.1.3 Isotopes of Hydrogen : If different

atoms of the same element have different

mass numbers they are called isotopes of each

other. (Refer to chapter 4). Hydrogen has three

isotopes with mass numbers 1, 2 and 3. They

all contain one proton and one electron but

different number of neutrons in the nucleus.

(see figure 8.1 and table 8.1)

p pn p n
n

Can you tell? 1 H 21H 13H
1
In which group should hydrogen be
placed ? In group 1 or group 17 ? Why ? Hydrogen Deuterium Tritium

Figure 8.1 Isotopes of hydrogen

110

Table 8.1 : Characteristics of isotopes of hydrogen

Name of the Symbol Atomic Atomic Neutron Abundance Stability
Isotope number Z mass number N
number A

Hydrogen H or 1 H or H-1 1 1 0 99.98 % Stable
1

Deuterium D or 2 H or H-2 1 2 1 0.015 % Stable
Tritium 1 1 3

T or 3 H or H-3 2 Trace Ratio
1 active

Do you know ? Stage 1 : Reaction of steam on hydrocarbon
or coke (C) at 1270 K temperature in presence
Tritium 3 H is a radioactive of nickel catalyst, gives water-gas, which is a
1 mixture of carbon monoxide and hydrogen.
nuclide with half life period 12.4 years

and emits low energy β particles.

8.1.4 Preparation of dihydrogen CH4(g) + H2O (g) 12N70i K CO(g) + 3H2(g)
}
Hydrogen can be prepared using many water-gas

methods. As water-gas is used for synthesis of CH3OH
and a number of hydrocarbons, it is also called
A. Laboratory methods

i. Dihydrogen is prepared in laboratory by the ‘syngas’. Production of syngas is also the first

action of dilute hydrochloric acid on zinc stage of gasification of coal.

granules. C(s) + H2O (g) 1270 K CO(g) + H2(g).
Sawdust, scrapwood, etc. can also be used in
Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)
ii. Dihydrogen can be prepared by the action

of aqueous solution of sodium hydroxide on place of carbon.

zinc. Stage 2 : Water-gas shift reaction: The carbon

Zn(s) + 2NaOH(aq) → Na2ZnO2(aq) + monoxide in the watergas is transformed
H2(g)
into carbondioxide by reacting with steam in
B. Industrial methods
presence of iron chromate as catalyst. This is
i. By electrolysis of pure water: Pure water is
called water-gas shift reaction.
a poor conductor of electricity. Therefore a
CO(g) + H2O(g) 673 K CO2(g) + H2(g)
dilute aqueous solution of acid or alkali is iron chromate
catalyst
used to prepare dihydrogen by electrolysis.
Stage 3 : Carbon dioxide is removed by
For example, electrolysis of dilute aqueous
scrubbing with sodium arsenite solution.
solution of sulphuric acid yields two volumes
Today major industrial production of
of hydrogen at cathode and one volume of
dihydrogen is from petrochemicals (∼ 77 %),
oxygen at anode. about 18 % from coal, about 4 % by electrolytic

2H2O (l) electrolysis akali 2H2↑ + O2↑ methods and about 1 % by other methods.
trace of acid or
8.1.5 Properties of dihydrogen
Pure dihydrogen (> 99.5% purity) gas is
A. Physical properties : Dihydrogen is
obtained by electrolysis of warm solution of
colourless, tasteless and odourless gas. It burns
barium hydroxide between nickel electrodes.
with a pale blue flame. It is a nonpolar water
ii. From carbon or hydrocarbon: Three
insoluble gas, lighter than air.
stages are involved in this industrial process

of preparation of dihydrogen.

111

B. Chemical properties : CuO(s) + H2(g) Cu(s) + H2O(l)

i. Reaction with metals : Dihydrogen Fe3O4 (s) + 4 H2(g) 3Fe(s) + 4H2O(l)

combines with all the reactive metals such as Pd2⊕(aq) + H2 (g) Pd(s) + 2H⊕(aq)

alkali metals, calcium, strontium and barium

at high tempereture, to form metal hydrides. b. Hydrogenation of unsaturated organic

For example : compound

2Na(s) + H2(g) 2NaH(s) The hydrogenation of unsaturated organic
compounds such as oil using nickel caralyst
(In this respect dihydrogen is similar to gives saturated organic compounds such as
solid fat (vanaspati Ghee).
halogens which also react with metals and

form metal halides.)

Just think C=C +H2 Ni CH - C H
(Saturated fat)
In the above chemical reaction (Vegetable oil)
which element does undergo oxidation
and which does undergo reduction ? 7.1.6 Uses of dihydrogen

Do you know ? i. Largest use of dihydrogen is in production
of ammonia.

The bond dissociation energy of ii. Dihydrogen is used in formation of
H-H bond is high, which is 436 kJ mol-1. vanaspati ghee by catalytic hydrogenation
Therefore reactions of dihydrogen take place of oils.
at high temperature and /or in the epresence
of catalyst. iii. Liquid dihydrogen is used as a rocket fuel.

ii. Reaction with dioxygen: Dihydrogen reacts iv. Dihydrogen is used in preparation of
with dioxygen in the presence of catalyst or by important organic compounds like
heating to form water. This reaction is highly methanol in bulk quantity.
exothermic.
2H2(g) + CO(g) Cobalt catalyst CH3OH(l)
2H2(g) + O2(g) catalyst 2H2O(l) ;
or heating v. Dihydrogen is used for preparation of
hydrogen chloride (HCl) and metal
∆H = -235 kJ mol-1 hydrides.

iii. Reaction with halogens: Dihydrogen Problem 8.1 : Justify the placement of
inflames with fluorine even at -2500 C in
dark, whereas it requires catalyst to react with hydrogen in the group of alkali metals
iodine. The vigour of reaction of dihydrogen
decreases with increasing atomic number of with the help of reaction with halogens.
halogen.
Solution :
H2 (g) + X2(g) → 2HX(g)
Hydrogen on reaction with halogen (X2)
(In this respect dihydrogen resembles alkali gives compounds with general formula
metals which also react with halogens to form
halides.) HX. For example : H2 + Cl2 2HCl

iv. Reducing nature of dihydrogen : similarly alkali metals (M) on reaction

a. Dihydrogen reduces oxides and ions of with halogens (X2) give compounds with
metals those are less reactive than iron, to the general formula MX.
corresponding metals at moderate temperature.
For example : For examples : 2Na + Cl2 2NaCl
Thus, H2 and alkali metal are monovalent
elements more electropositive than

halogens. This similarily justifies the

place of hydrogen in the group 1.

112

Can you recall? Group 2 of the periodic table consists of
elements : beryllium, magnesium, calcium,

1. What is the name of the family of strontium, barium and radium. These
reactive metals having valancy one? elements are collectively called alkaline earth
2. What is the name of the family of metals because they occur as minerals in
reactive metals having valancy two? rocks. The elements magnesium and calcium
are found abundantly in earth crust but
8.2 Alkali metals and alkaline earth metals radium is not easy to find. Radium is one of
the first two radioactive elements discovered
8.2.1 Introduction : The elements of the groups by Madame Curie.
1 and 2 are placed on the left in the periodic 8.2.2 Electronic configuration of elements of
table. Here the last electron enters into ‘ns’ group 1 and group 2
subshell. Thus they belong to the s-block of the
periodic table. Group 1 of the periodic table The general outer electronic configuration
consists of the elements: hydrogen, lithium, of the group 1 elements is ns1 and that of the
sodium, potassium, rubidium, caesium and group 2 elements is ns2. The loosely held
francium. These elements except hydrogen s-electrons in valence shell of these elements
are collectively called alkali metals. We have can be easily removed to form metal ions.
already looked at hydrogen in the section 8.1. Hence these elements are never found in free
Two elements of Group 1, namely, sodium state in nature. Tables 8.2 and 8.3 show the
and potassium are the sixth and seventh most electronic configurations of the elements of
abundant elements in earth crust but francium

does not occur appreciably in nature because group 1 and group 2 elements, respectively.

it is radioactive and has short half-life period.

Table 8.2 : Electronic configuration of group 1 elements

Name Symbol Atomic Condensed electronic Electronic Configuration.
number configuration

Hydrogen H 1 1s1 1s1

Lithium Li 3 [He] 2s1 1s2 2s1

Sodium Na 11 [Ne] 3s1 1s2 2s2 2p6 3s1

Potassium K 19 [Ar] 4s1 1s2 2s2 2p6 3s2 3p6 4s1
Rubidium Rb 37 [Kr] 5s1
Caesium Cs 55 [Xe] 6s1 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10
5p6 6s1

Francium Fr 87 [Rn]7s1 1s2 2s2 2p6 3s2 3p6 4s23d10 4p6 5s2 4d10
5p6 6s2 4f14 5d10 6p6 7s1

Table 8.3 : Electronic configuration of group 2 elements

Name Symbol Atomic Condense electronic Electronic Configuration.
Beryllium
Magnesium number configuration
Calcium
Strontium Be 4 [He] 2s2 1s2 2s2

Barium Mg 12 [Ne] 3s2 1s2 2s2 2p6 3s2

Radium Ca 20 [Ar] 4s2 1s2 2s2 2p6 3s2 3p6 4s2

St 38 [Kr] 5s2 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10
Ba 56 [Xe] 6s2
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2
Ra 88 [Rn] 7s2 4d10 5p6 6s2

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2
5p6 6s2 4f14 5d10 6p6 7s2

113

8.2.3 Trends in atomic and physical properties and soft, but harder than the alkali metals. They
of elements of group 1 and group 2 are also strongly electropositive in nature.
But comparatively less electropositive than
All the alkali metals are silvery white the alkali metals. Some atomic and physical
and soft. Due to their large atomic size these properties of the alkali metals and the alkaline
elements have low density. They are the most earth metals are listed in the tables 8.4 and 8.5
electropositive elements. The alkaline earth respectively.
are metals also in general silvery white lustrous

Table 8.4 : Physical properties of group 1 elements (except hydrogen)

Symbol Atomic Ionic Density Ionization Electro- Melting Abundance Standard
radius radius (g/cm-3) enthalpy negativity point (K) in the reduction
Li (pm) (pm) (kJ mol-1) potential
Na lithosphere
K Eo (V)
Rb 152 76 0.54 520 1.0 454 18 ppm -3.04
Cs 186 -2.714
Fr 227 102 0.97 496 0.9 371 2.27 % -2.925
248 -2.930
Symbol 265 138 0.86 419 0.8 336 1.84 % -2.927

Be - 152 1.53 403 0.8 312 78.12 ppm -
Mg
Ca Atomic 167 1.90 376 0.7 302 2.6 ppm Standard
Sr radius reduction
Ba (pm) (180) - ∼375 - - -10-18 ppm potntial
Ra
111 Table 8.5 : Physical properties of group 2 elements E0 (V)
160
197 Ionic Density Ionization Electro- Melting Abundance -1.97
215 radius (g/cm-3) enthalpy negativity point in the -2.36
222 (pm) ( kJ mol-1) (K) -2.84
1st 2nd lithosphere -2.89
- -2.92
31 1.84 899 1757 1.5 1560 2 ppm -2.92

72 1.74 737 1450 1.2 924 2.76 %
1124 4.6 %
100 1.55 590 1145 1.0 1062 384 ppm
1002 390 ppm
118 2.63 549 1064 1.0 973 10-6 ppm

135 3.59 503 965 0.9

148 (5.5) 509 979 -

Unipositive ions of all the elements of both the groups. Ionization enthalpies and

group 1 have inert gas configuration. Thus electronegativities decrease down both the

they have no unpaired electron and their groups. The elements of both these groups, in

compounds are diamagnetic and colourless. general, have high negative values of standard

The divalent ions of group 2 elements also reduction potentials.

have inert gas configuration with no unpaired 8.2.4 Chemical properties of elements of

electron, and therefore their compounds are group 1 and group 2

also diamagnetic and colourless. Lithium and The alkali metals and alkaline earth

beryllium differ from the rest of the elements metals are very reactive in nature. As a result of

of the groups 1 and 2, respectively because of this they are always found in combined state.

their extremely small size and comparatively Their reactivity is due to their low ionization

high electronegativity. ehthalpy values in general. The reactivity

The physical properties of group 1 and of these metals increases with increasing

group 2 elements show reasonable regularity atomic radius and corresponding lowering of

in the periodic trends. Thus the atomic and lonization enthalpy down the groups 1 and 2

ionic radii and densities increase down can be noticed.

114

Problem 8.2: Problem 8.4:
Sodium forms ionic compounds having What is the oxidation state of Na in
formulae NaCl, NaH and Na2CO3. Explain Na2O2?
Solution: Let us rewrite the formulae Solution: The peroxide species is
of the compounds of sodium showing represented as O22−. Any compound is
charges on the concerned cation and electrically neutral. Therefore oxidation
basic anion. state of each Na is (+2/2 = +1) in
Na⊕Cl , Na⊕H , 2Na⊕ CO32 Na2O2.
It is seen that in all these compounds
Na carries one positive charge. The Na⊕ Problem 8.5 :
is formed from Na atom by losing one The atomic radii of Na, K and Mg
electron. Na → Na⊕+ e are 186, 227 and 160 pm, respectively.
This happens because the electronic Explain the differences.
configuration of Na is [Ne]3s1. There is Solution : Na and K both belong to the
only one electron in the valence shell of group 1. K has larger valence shell than
Na. It can easily be lost as the ionization Na. Therefore the atomic radius of K
enthalpy is low. And Na⊕ ion so formed is larger than that of Na. Na and Mg
is stable as it has stable electronic belong to the same period. Therefore both
configuration of the inert gas Ne. have the same valence shell. But the
nuclear charge of Mg is larger than that
Problem 8.3 : Explain the observed of Na. Therefore, the valence electrons of
Mg are held more tightly and its atomic
values 496 kJ/mol and 737 kJ/mol of the radius is smaller than that of Na.

first ionization enthalpies of Na and Mg, The elements of group 1 and group 2 both

respectively. being s-block elements, show similarily in

Solution : The electronic configuration their chemical properties. The differences are

of Na is [Ne]3s1 and that of Mg is due to variation in the atomic radii, ionization

[Ne]3s2. During the first iouization only enthalpies and valencies.

one electron is removed from a neutral i. Reaction with oxygen/air

atom. Na⊕ + e Group 1 - All the elements of group 1 rapidly
Na 1st ionisation Mg⊕ + e
Mg 1st ionisation lose their luster in air due to formation of a

The resulting Na is isoelectronic with Ne, layer of oxide, on peroxide and in some cases

and therefore, is stable. Thats why Na superoxide by reaction with oxygen in air.

has low value of 1st ionization enthalpy. 2Li + O2 2LiO

However, to form the unipositive Mg⊕ (Lithium oxide)

ion energy is required to unpair these 2Na + O2 Na2O2
(Sodium peroxide)
electrons in the valency shell and also
K + O2 KO2
remove one electron to form the ion
(Potassium superoxide)
having electronic configuration [Ne] 3s1.

It is not as stable as Na since it does not Do you know ?

correspond to any inert gas. Therefore the

first ionization enthalpy of Mg is higher • The reaction of Na and K with oxygen
is highly exothermic and these metals
than that of Na. catch fire when exposed to air.

115

• Potassium superoxide has ability to Group 2 : All the metals of group 2, except
absorb carbon dioxide and give out
oxygen at the same time: beryllium, when heated with hydrogen form

4KO2 + 2CO2 → 2K2CO3 + 3O2↑ MH2 type hydrides. MH2
• This property of KO2 has been made M + H2

use of in breathing equipment used Problem 8.6 : NaCl is an ionic
for mountaineers and in submarines compound but LiCl has some covalent
and space. character, explain.
Solution: Li⊕ ion has very small size,
The oxides of group 1 metals are strongly therefore the charge density on Li⊕ is
high. Therefore it has high tendency
basic in nature. They dissolve in water to distort the electron cloud around the
negatively charged large chloride ion.
forming aqueous solutions of strong alkali. This results in partial covalent character
of the LiCl bond. Na⊕ ion cannot distort
For example the electron cloud of Cl due to the
bigger size of Na⊕ compared to Li⊕.
2LiO(s) + H2O(l) → 2LiOH(aq)
Group 2 : These metals are protected from

air oxidation by an oxide film formed on their

surface.

All the elements of group 2 burn when ignited iv. Reaction with Halogens

in air forming MO type oxides. The product is Group 1 : All the alkali metals react vigorously

a mixture of oxide and nitride. with halogens to produce their ionic halide

2Mg + O2 → 2MgO salts.
3Mg + N2 → Mg3N2
Further heating of the oxide in air results in 2M + X2 2M+ X-

Group 2: All the alkaline earth metals

formation of peroxide. combine with halogens at high temperature to

ii. Reaction with water form halides.

Group 1 : Lithium, sodium and potassium all M + X2 MX2
v. Reducing nature
float on water due to hydrogen bubbles released

on reaction with water. Lithium reacts slowly Group 1 : The reducing power of an

but sodium and potassium react vigrously element is measured in terms of standard

with water. Due to highly exothermic reaction electrode potential (E0) corresponding to the

sodium and potassium catch fire when put in transformation M⊕(aq) + e → M (s). All the

water. alkali metals have high negative values of

2Na + 2H2O → 2Na OH + H2↑ E0 indicative of their strong reducing nature,
Group 2 : The elements of group 2 react
lithium is the most powerful and sodium is the

with water to form metal hydroxide and least powerful in the group. (see Table 8.4)

hydrogen. Beryllium does not react with Group 2 : All the alkaline earth metals have

water. Magnesium decomposes hot water, high negative values of standard reduction

other elements react with cold water forming potential (E0). (see Table 8.5), and are strong

metal hydroxide M(OH)2 and hydrogen gas. reducing agents. However their reducing power
Ca + 2H2O → Ca(OH)2↓ + H2↑
is less than those of alkali metals.
iii. Reaction with Hydrogen
vi Solution in liquid ammonia

Group 1 : Alkali metals react with hydrogen Group 1 : The alkali metals are soluble in

at high tempereture to form the corresponding liquid ammonia giving deep blue coloured

metal hydrides. solutions which show electrical conductivity.
2M + H2 673 K
2M+ H- M + (x + y)NH3 → [M(NH3)x]⊕ + [e(NH3)y]

116

The ammoniated electron is responsible remaining alkali metals and resembles with

for the deep blue colour of these solutions. magnesium, the second alkaline earth metal.

These solutions are paramagnetic and on Likewise beryllium shows many differences

standing liberate hydrogen slowly, resulting in with remaining alkaline earth metals and

formation of the metal amide. The blue colour shows similarity with aluminium, the second

changes to bronze and the solution becomes element of the next main group (group 13).

diamagnetic. The relative placement of these elements with

M⊕(am) + e (am) + NH3 (l) → MNH2(am) similar properties in the periodic table appears
+ H2(g) to be across a diagonal (see. Table 8.6) and is
called diagonal relationship.
(Here (am) denotes solution in ammonia.)

Group 2 : Similar to alkali metals the alkaline Table 8.6 : Diagonal relationship

earth metals are also soluble in liquid ammonia Main Group 1 2 13
which give deep blue black coloured solutions. Period

M + (x + 2y) NH3 →[M(NH3)x]2⊕+ 2[e(NH3)y] 2 Li Be B
8.2.5 Diagonal Relationship : It is expected
that elements belonging to the same group 3 Na Mg Al

exhibit similarity and gradation in their The table 8.7 shows some properties of

properties. The first alkali metal lithium and lithium and magnesium which elucidate their

the first alkaline earth metal beryllium do not diagonal relationship.

fulfil this expectation. Thus, lithium shows

many differences when compared with the

Table 8.7 : Ressemblence between Li and Mg

Criterion Products of Products of thermal Property of Formula of
decomposition of carbonate chloride crystalline
reaction with air M2CO3/MCO3 ∆ ? chloride
M air ?
Element

Li Li2O + Li3N Li2O + CO2 deliquescent LiCl.2H2O

Mg MgO + Mg3 N2 MgO + CO2 deliquescent MgCl2.8H2O

Group 1 M2O/M2O2/MO No reaction Not MCl
(except Li) deliquescent

Group 2 MO + M3N2 MO + CO2 deliquescent MCl2.xH2O

In table 8.8 some properties of Be and Al are shown which indicate the diagonal relationship

Table 8.8 : Resemblence between Be and Al

Criterion Properties of chloride
Element
Nature of bonding Whether Lewis Solubility Properties of oxide
acid in organic Acidic/Basic/Amphoteric

solvent

Covalent
chain structure with
Cl bridges BeLCelw2 iiss strong Amphoteric
Cl Cl acid
Be Soluble BeO+2HCl BNeCa2lB2+eHO22+H2O
Be Be Be BeO+2NaOH
Cl Cl

117

Covalent
dimer with Cl
bridges Amphoteric

Al Cl Cl Cl AlCl3 is strong Soluble AAll22OO33++26NHaOClH 2AlCl3+3H2O
Al Al Lewis acid 2NaAlO2+H2O

Cl Cl Cl

Group 2 Ionic Not Insoluble Basic
Group 13 Covalent Lewis acid Soluble MMOO++HNCalO→H→MCNl2o+reHa2ction

Lewis acid Amphoteric

The diagonal relationship between the v. Radium is used in radiotherapy for cancer
treatment.
elemental pairs belonging to different groups
8.2.7. Biological importance of elements of
and periods is due to the similarity in some of group 1 and group 2
Group 1
their atomic properties. Thus atomic and ionic i. Sodium ion is present as the largest

radii of Li and Mg are very similar. (see the supply in all extracellular fluids. These
fluids provide medium for transportting
table 8.4). In the case of Be and Al the charge nutrients to the cells.
ii. The concentration of sodium ion in
to radius ratio of their ions is very similar. (Be extracellular fluids regulates the flow of
: 2 and Al : 3 ) water across the membrane.
iii. Sodium ions participate in the
31 53.55 transmission of nerve signals.
8.2.6 Uses of elements of group 1 and group 2 iv. Potassium ions are the most abundant
Group 1 ions within cells. These are required for
i. Lithium metal is used in long-life batteries maximum efficiency in the synthesis of
proteins and also in oxidation of glucose.
used in digital watches, calculators and
computers. Group 2
ii. Liquid sodium has been used for heat i. Mg2⊕ ions are important part of
transfer in nuclear power station.
iii. Potassium chloride is used as a fertilzer. chlorophyll in green plants.
iv. Potassium is used in manufacturing ii. Mg2⊕ ions play an important role in the
potassium superoxide (KO2) for oxygen
generation. It is good absorbent of carbon breakge of glucose and fat molecules, in
dioxide. synthesis of proteins with enzymes, and
v. Caesium is used in photoelectric cells. in regulation of chlolesterol level.
Group 2 iii. Ca2⊕ ions are important for bones and
i. Beryllium is used as a moderator in teeth in the form of apatite [Ca3(PO4)2]
nuclear reactors. iv. Ca2⊕ ions play important role in blood
ii. Alloy of magnesium and aluminium is clotting.
widely used as structural material and in v. Ca2⊕ ions are required for contraction and
aircrafts. stretching of muscles.
iii. Calcium ions are important ingredient in vi. Ca2⊕ ions are also required to maintain
biological system, essential for healthy the regular beating of heart.
growth of bones and teeth.
iv. Barium sulphate is used in medicine as
barium meal for intestinal x-ray.

118

Problem 8.7 : Magnesium strip slowly In the second stage the separated crystals
tarnishes on keeping in air but metallic
calcium is readily attacked by air. Explain. of sodium bicarbonate are heated to obtain
Solution: Mg and Ca belong to group 2,
but periods 2 and 3, respectively. During the sodium carbonate.
reaction with air the metallic Mg and Ca
lose their valence electrons. The tendency 2 NaHCO3(s) ∆ Na2CO3(s) + H2O(g)
to lose valence electron is the metallic + CO2(g)
character, which increases down the group.
Thus, calcium has higher metallic character, In this process the recovery of ammonia is
greater tendency to lose valence electron and
lower ionization enthalpy than magnesium. done by treating the solution of NH4Cl obtained
Therefore Mg reacts slowly with air, with slaked lime, Ca(OH)2. The byproduct of
forming a thin film of oxide, resulting into this reaction is calcium chloride.
tarnishing, whereas Ca reacts readily at
room temperature with oxygen and nitrogen 2 NH4Cl(aq) + Ca(OH)2(s) 2 NH3(g)
in the air. + CaCl2(aq) + H2O(l)

Do you know ?

Potassium carbonate can not be
prepared by Solvay process because
potassium hydrogen carbonate is highly
water soluble and cannot be precipitated
by reaction with potassium choride.

8.3 Some important compounds of elements Properties : Sodium carborate (washing soda)

of s-block is a white crystalline solid having the formula

In this section we consider five important Na2CO3, 10H2O. It is highly soluble in water.
On heating the decahyadrate loses water
compounds of s-block elements with reference
molecules to form monohydrate. On heating
to their prepartion, properties and uses.
above 373 K temperature monohydrate further
8.3.1. Sodium Carbonate (washing soda)
loses water and changes into white anhydrous
Na2CO3. 10H2O
Prepartion: Sodium Carbonate is powder called soda-ash.
Na2CO3.10H2O(s) 373 K Na2CO3.H2O(s) +
commercially prepared by Solvay process. 9H2O(g)
Na2CO3(s)+H2O(g)
In the first stage of process CO2 is passed Na2CO3.H2O(s) > 373 K
into a concentrated solution of NaCl which Aqueous solution of sodium carbonate is

is saturated with NH3. Crystals of sodium alkaline because of its hydrolysis by the
bicarbonate separate as a result of the
following reaction:
following reactions.
Na2CO3 + H2O → NaHCO3 + NaOH
Reactions in the first stage: Uses

2 NH3(aq) + H2O + CO2 (g) (NH4)2 CO3(aq) i. The alkaline properties of sodium carbonate
(NH4)2CO3(aq) + H2O + CO2(aq)
are responsible for emulsifying effect on
2 NH4HCO3(aq)
NH4HCO3(aq) + NaCl(aq) NH4Cl(aq) grease and dirt. It is used as cleaning material.

+ NaHCO3(s) ii. It is used to make hard water soft (as a water
Sodium bicarbonate has low solubility,
softener), as it precipitates out the soluble

and therefore its crystals precipitate out calcium and magnesium salts in hard water as

which are formed as a result of the double carbonates. For example : Ca(HCO3)2(aq) +
Na2CO3(aq) → CaCO3(s)
decomposition reaction between ammonium
+ 2NaHCO3(aq)
bicarbonate and sodium chloride. iii. It is used for commercial production of

soap and caustic soda.

iv. It is an important laboratory reagent.

119

8.3.2 Sodium hydroxide (caustic soda) Properties
NaOH
Calcium carbonate is soft, light, white powder.
Preparation : Sodium hydroxide is It is practically water insoluble. On heating to
commercially obtained by the electrolysis of 1200 K calcium carbonate decomposes into
saturated aqueous solution of sodium chloride. calcium oxide and carbon dioxide.
Brine solution is subjected to electrolysis CaCO3(s) 1200 K CaO(s) + CO2↑
in Castner-Kellner cell. Mercury is used as It reacts with dilute acids to give the

cathode and carbon rod as anode. Metallic corresponding calcium salt and carbon
sodium liberated at the cathode forms sodium dioxide.
amalgam. Chlorine gas is evolved at the
anode. CaCO3 + 2HCl → CaCl2 + CO2 ↑+ H2O
Cathode reaction: Na⊕+e Hg Na-amalgam CaCO3 + H2SO4 → CaSO4 + CO2↑+ H2O
Uses:

Anode reaction: Cl 1/2 Cl2 + e i. Calcium carbonate in the form of marble

Sodium hydroxide is obtained by treating is used as building material.

sodium amalgam with water, when hydrogen ii. Calcium carbonate is used in the

gas is liberated. manufacture of quicklime (CaO) which is

2Na -Hg + 2H2O 2NaOH + 2Hg + H2 the major ingredient of cement.

Properties : Sodium hydroxide is a white iii. A mixture of CaCO3 and MgCO3 is used as
flux in the extraction of metals from ores.
deliquescent solid, having melting point 591 K.
iv. It is required for the manufacture of high
It is highly water soluble and gives a strongly
quality paper.
alkaline solution. The surface of the solution
v. It is an important ingredient in toothpaste,
absorbs atmospheric CO2 to form Na2CO3.
Uses chewing gum, dietary suppliments of

i. Sodium hydroxide is used in purification of calcium and filler in cosmetics.

bauxite (the aluminium ore). 8.3.4 Hydrogen peroxide (H2O2) : Hydrogen
peroxide is a low cost, clean and mild oxidising
ii. It is used in commercial production of soap,
agent. A 30% aqueous solution hydrogen
paper, artificial silk and many chemicals.
peroxide is commercially available.
iii. It is used for mercerising cotton fabrics.
Preparation :
iv. It is used in petroleum refining.
i. Hydrated barium peroxide is treated with
v. It is an important laboratory reagent.
ice cold dilute sulfuric acid. The precipitate of
8.3.3 Calcium Carbonate (CaCO3)
Calcium carbonate is found in nature as chalk, barium sulphate formed is filtered off to get

lime stone, marble. hydrogen peroxide solution.

Preparation BaO2.8H2O(s) + H2SO4(aq) low temp BaSO4↓
273 K
+ H2O2(aq) + 8H2O(l)
i. When carbon dioxide is bubbled through
ii. Small quantity of sodium peroxide is added
solution of calcium hydroxide (slaked lime)
to ice-cold solution of dilute sulfuric acid
water insoluble solid calcium carbonate is
with strirring gives hydrogen peroxide (Merck
formed.
process).
Ca(OH)2(aq) + CO2(g) CaCO3(s) + H2O(l)
Excess carbon dioxide transforms the Na2O2 (aq) + H2SO4 (aq) 273 K H2O2 (aq)
+ Na2SO4 (aq)
precipitate of CaCO3 into water soluble calcium
bicarbonate and therefore has to be avoided. iii. A 50 % solution of sulfuric acid is

ii. When solution of calcium chloride is added subjected to an electrolytic oxidation to form
to a solution of sodium carbonate, calcium peroxydisulfuric acid at anode.
carbonate is formed as precipitate.
2HSO4 Electrolysis H2S2O8 + 2e
CaCl2(aq)+Na2CO3(aq) CaCO3(s)+ 2NaCl(aq)

120

Hydrolysis of the peroxydisulfuric acid yields excess chlorine from fabrics which have been

hydrogen peroxide. bleached by chlorine.

HO-SO2-O-O-SO2-OH + 2H2O → 2H2SO4 iii. Now a days it is also used in environmental
+ H2O2 chemistry for pollution control, restoration of
aerobic condition to sewage water.
This method can be extended to laboratory

preparation of D2O2 Promblem 8.8 : Calculate % (by mass) of
iv. Industrially hydrogen peroxide is prepared
a H2O2 solution which is 45.4 volume.
by air-oxidation of 2-ethylanthraquinol. Solution: 45.4 Volume H2O2 solution
means 1L of this solution will give 45.4 L
The 2-ethylanthraquinol is regenerated

by catalytic hydrogenation of O2 at STP

2-ethylanthraquinone. O2 2 H2O2 → 2 H2O + O2
air 22.7 L at STP
2-ethylanthraquinol H2O2 (2×34) g

+ 2 - ethylanthraquinone Thus, 22.7 L O2 at STP is produced by 68
g H2O2
OH O2 O
air C2H5 + H2O2 ∴45.4 L O2 at STP is produced by 68× 45.4
= 136 g H2O2 22.7
C2H5 O
OH H2/Pd/∆
∴Strength this H2O2 solution = 136 g/L
Properties

i. Pure H2O2 is a very pale blue coloured = 136 g H2O2 =13.6% (by mass)
liquid, having b.p. 272.4 K. 1000 g water

ii. H2O2 is miscible in water and forms a 8.3.5 Lithium aluminium hydride (LiAlH4)
hydrate (H2O2. H2O) Lithium aluminium hydride is commonly
abbreviated as LAH. It has chemical formula
iii. Strength of aqueous solution of H2O2 LiAlH4.
is expressed in ‘volume’ units. The Prepartion : Lithium hydride is treated with
commercially marketed 30% (by aluminium chloride to give lithium aluminium
mass) solution of H2O2 has volume hydride
strength of 100 volume. It means
that 1 mL of 30% solution of H2O2 will 4LiH + AlCl3 → LiAlH4 + 3LiCl
give 100 mL oxygen at STP. Properties: Lithium aluminium hydride is a
colourless solid. It reacts violently with water
iv. H2O2 acts as a mild oxidising as well as and even atmospheric moisture.
reducing agent. Uses
(i) LAH is a source of hydride and therefore
a. Oxidising action of H2O2 in acidic used as reducing agent in organic synthesis.
medium
O
2Fe2⊕(aq) + 2H⊕(aq) + H2O2(aq) → 2Fe3⊕(aq)
+ 2H2O(l)

b. Reducing action of H2O2 in acidic
medium

2 MnO4 + 6H⊕ +5H2O2 → 2Mn2⊕ + 8H2O R C OR' LiAlH4 R OH + R' - OH
+ 5O2 dry ether

Uses O

i. Hydrogen peroxide is used as mouthwash, R C LiAlH4 R OH + H2O
germicide, mild antiseptic, preservative for dry ether
milk and wine and bleaching agent for soft
OH

materials due to its mild oxidising property. (ii) LAH is useful to prepare PH3(phosphine)
ii. Hydrogen peroxide, due to its reducing 4 PCl3 + 3LiAlH4 → 4 PH3 + AlCl3 + Li Cl
property, is used as an antichlor to remove

121

Exercises

1. Explain the following 4. Name the following
A. Hydrogen shows similarity with alkali
metals as well as halogens. A. Alkali metal with smallest atom.
B. Standard reduction potential of alkali
metals have high negative values. B. The most abundant element in the
C. Alkaline earth metals have low values
of electronegativity; which decrease universe.
down the group.
D. Sodium dissolves in liquid ammonia to C. Radioactive alkali metal.
form a solution which shows electrical
conductivity. D. Ions having high concentration in cell
E. BeCl2 is covalent while MgCl2 is ionic.
F. Lithium floats an water while sodium sap.
floats and catches fire when put in
water. E. A compound having hydrogen,

2. Write balanced chemical equations for aluminium and lithium as its constituent
the following.
A. CO2 is passed into concentrated elements.
solution of NaCl, which is saturated
with NH3. 5. Choose the correct option.
B. A 50% solution of sulphuric acid is
subjected to electrolyte oxidation and A. The unstable isotope of hydrogen is .....
the product is hydrolysed.
C. Magnesium is heated in air. a. H-1 b. H-2
D. Beryllium oxide is treated separately
with aqueous HCl and aqueous NaOH c. H-3 d. H-4
solutions.
B. Identify the odd one.
3. Answer the following questions
A. Describe the diagonal relationship a. Rb b. Ra
between Li and Mg with the help of
two illustrative properties. c. Sr d. Be
B. Describe the industrial production of
dihydrogen from steam. Also write the C. Which of the following is Lewis acid ?
chemical reaction involved.
C. A water sample, which did not give a. BaCl2 b. KCl
lather with soap, was found to contain
Ca(HCO3)2 and Mg(HCO3)2. Which c. BeCl2 d. LiCl
chemical will make this water give
lather with soap? Explain with the help D. What happens when crystalline Na2CO3
of chemical reactions. is heated ?
D. Name the isotopes of hydrogen.
Write their atomic composition a. releases CO2
schematically and explain which of b. loses H2O
these is radioactive ? c. decomposes into NaHCO3
d. colour changes.

Activity :

1. Collect the information of preparation of
dihydrogen and make a chart.

2. Find out the s block elements compounds
importance/uses.

122

9. Elements of Groups 13, 14 and 15

Can you recall? and lead (82Pb) form the group 14 called the
carbon family. The elements nitrogen (7N),
If the valence shell electronic phosphorous (15P), arsenic (33As), antimony
configuration of an element is 3s2 3p1 in (51Sb) and bismuth (83Bi) belong to group 15 of
which block of periodic table is it placed ? the periodic table called the nitrogen family.
9.2 Electronic configuration of elements of
9.1 Introduction : You have learnt in groups 13, 14 and 15
Chapter 7 that in the p-block elements
the diffrentiating electron (the last filling The general outer electronic configuration
electron) enters the outermost p orbital. You

also know that maximum six electrons can of the group 13 elements is ns2 np1, those of the

be accommodated in p-subshell (or three p group 14 elements is ns2np2 while the group 15
orbitals). This gives rise to six groups, group elements are shown as ns2np3. These electronic
13 to 18, in the p-block. The p-block elements configurations differ from their nearest inert
show greater variation in the properties than gas by 3 or 4 electrons. These elements do
's' block, which you learnt in the previous not occur in free monoatomic state and found
chapter. In this chapter you are going to study as compounds with other elements or as
the elements of the groups 13, 14 and 15 in polyatomic molecules (such as N2, P4, C60) or
some details. polyatomic covalent arrays (such as graphite,

The elements boron (5B), aluminium diamond). Table 9.1 shows the condensed
(13Al), gallium (31Ga), indium (49In) and electronic configuration of the elements of
thallium (81Tl) constitute the group 13, called group 13, group 14 and group 15.
the boron family. The elements carbon (6C),
silicon (14Si), germanium (32Ge), tin (50Sn)

Table 9.1 : Condensed electronic configuration of elements of groups 13, 14 and 15

Group 13 (Boron family) Group 14 (Carbon family) Group 15 (Nitrogen family)

Element Condesed electronic Element Condesed electronic Element Condesed electronic

configuration configuration configuration

5B [He]2s22p1 6C [He]2s22p2 7N [He]2s22p3

13Al [Ne]3s23p1 14Si [Ne]3s23p2 15P [Ne]3s23p3

31Ga [Ar]3d104s24p1 32Ge [Ar]3d104s24p2 33As [Ar]3d104s24p3

49In [Kr]4d105s25p1 50Sn [Kr]4d105s25p2 51Sb [Kr]4d105s25p3

81Tl [Xe]4f145d106s26p1 82Pb [Xe]4f145d106s26p2 83Bi [Xe]4f145d106s26p3

Problem 9.1 : Atomic numbers of the group 13 elements are in the order B < Al < Ga < In <
Tl. Arrange these elements in increasing order of ionic radii of M3⊕.
Solution : The given elements are in an increasing order of atomic number. As we go down
the group 13, their general outer electronic configuration is ns2np1. M3⊕ is formed by removal
of three electrons from the outermost shell 'n'. In the M3⊕ the 'n-1' shell becomes the outermost.
Size of the added 'n-1' shell increases down the group. Therefore the ionic radii of M3⊕ also
increase down the group as follows :
B3⊕ < Al3⊕ < Ga3⊕ < In3⊕ < Tl3⊕

123

Problems 9.2 Why the atomic radius of 9.3 Trends in atomic and physical properties
Gallium is less than that of aluminium ? of elements of groups 13, 14 and 15. : The
Solution : Atomic radius increases down elements of group 13 show a wide variation in
the group due to added new shell. 'Al' does properties. In group 13 boron is a metalloid.It
not have ‘d’ electrons. As we go from Al is glossy and hard solid like metal but a poor
down to 'Ga' the nuclear charge increase electrical conductor (like nonmetals). The
by 18 units. Out of the 18 electrons added, other elements in this group are fairly reactive
10 electrons are in the inner 3d subshell. metals. Aluminium is the third most abundant
'd' Electrons offer poor shielding effect. element in the earth’s crust. Some physical
Therefore, the effects of attraction due to and atomic propeties of group 13 elements are
increased nuclear charge is experienced given in Table 9.2
prominently by the outer electrons of 'Ga'
and thus its atomic radius becomes smaller
than that of 'Al'.

Table 9.2 : Physical properties of elements of group 13

Element Atomic Atomic Atomic Ionic radius Ionization enthalpy Electronegativity Density Melting Boiling
B number point point
mass radius (pm) (kJ mol-1) (g/cm3) (K) (K)
5 24.53 3923
(pm) M3⊕ 1st 2nd 3rd

10.81 88 27 801 2427 3659 2.0 2.35

Al 13 26.98 143 53.5 577 1816 2744 1.5 2.70 933 2740

Ga 31 69.72 135 62.0 579 1979 2962 1.6 5.90 303 2676
558 1820 2704 1.7 7.31 430 2353
In 49 114.82 167 80.0 589 1971 2877 1.8 11.85 576 1730

Tl 81 204.38 170 88.5

Problem 9.3 : The values of the first ionization enthalpy of Al, Si and P are 577, 786 and
1012 kJmol-1 respectively. Explain the observed trend.
Solution : The trend shows increasing first ionization enthalpy from Al to Si to P. Al, Si and P
belong to 13 period in the periodic table. They have same valence shell. Due to the increased
nuclear charge electrons in the valence shell are more tightly held by the nucleus as we go
from Al to Si to P. Therefore more energy is required to remove an electron from its outermost
shell.

Table 9.3 enlists atomic and physical which is brittle like nonmetal, it is hard and
properties of the elements of carbon family has metallic luster. Germanium is also brittle
(group 14). In this group all the three but hard and lustrous metalloid. Tin or lead
traditional types of elements are present. down the group are corrosion resistant and
Carbon is a nonmetal, silicon is a metalloid moderately reactive.

Table 9.3 : Some atomic and physical properties of group 14 elements

Element Atomic Atomic Atomic Ionic radius Ionization enthalpy Electro- Density Melting Boiling
number mass radius (pm) (kJ mol-1) negativity (g/cm3) point point
(pm) (K)
M3⊕ 1st 2nd 3rd 4th (K)

C 6 12.01 77 1086 2352 4620 6220 2.5 3.51 4373

Si 14 28.09 118 40 786 1577 3228 4354 1.8 2.34 1693 3550

Ge 32 72.60 122 53 761 1537 3300 4400 1.8 5.32 1218 3123

Sn 50 118.71 140 69 708 1411 2942 3929 1.8 7.26 505 2896

Pb 82 207.2 146 78 715 1450 3081 4082 1.9 11.34 600 2024

124

Table 9.4 : Some atomic and physical properties of group 15 elements

Element Atomic Atomic Atomic radius Ionic Ionization enthalpy Electronegativity Density Melting Boiling

number mass (pm) radius (kJ mol-1) (g/cm3) point point

(pm) 1st 2nd 3rd (K) (K)

N 7 14.01 70 171 (M3 ) 1402 2856 4577 3.0 0.879 6.3 77.2

P 15 30.97 110 212 (M3 ) 1012 1903 2910 2.1 1.823 317 554

As 33 74.92 121 222(M3 ) 947 1798 2736 2.0 5.778 1089 -

Sb 51 121.75 141 76 (M3⊕) 834 1595 2443 1.9 6.697 904 1860

Bi 83 208.98 148 103(M3⊕) 703 1610 2466 1.9 9.808 544 1837

Group 15 is the nitrogen family. Table valence electrons (sum of s- electrons and

9.4 shows atomic and physical properties p-electrons). This is sometimes called the

of the elements of group 15. This group also group oxidation state. In boron, carbon and

include the three traditional types of elements: nitrogen families the group oxidation state is

The gaseous nitrogen and brittle phosphorous the most stable oxidation state for the lighter

are nonmetals. Arsenic and antimony are elements. Besides, the elements of groups

metalloids while bismuth is moderately 13, 14 and 15 exhibit other oxidation states

reactive metal. which are lower than the group oxidation

9.4 Chemical properties of the elements of state by two units. The lower oxidation states

the groups 13,14 and 15 become increasingly stable as we move down

9.4.1 Oxidation state : Oxidation state is the to heavier elements in the groups. Table 9.5

primary chemical property of all elements. shows various oxidation states exhibited by

The highest oxidation state exhibited by the elements belonging to these groups.

p-block elements is equal to total number of

Remember

The increased stability of the oxidation state lowered by 2 units than the group oxidation
state in heavier p-block elements is due to inert pair effect. In these elements, the two s-
electrons are involved less readily in chemical reactions. This is because the s-electrons of
valence shell in p-block elements experience poor shielding than valence p- electron, due to ten
inner d-electrons.

Table 9.5 : Oxidation states of the elements of groups 13, 14 and 15

Group 13 14 15

Outer electronic ns2 np1 ns2 np2 ns2 np3
configuration
+3 +4 +5
Group oxidation state +1 +2, -4 +3, -3
Other oxidation
states BF3, AlCl3, GaCl3, CH4, CO2, CCl4, SiCl4, N2O5, NH3, NF3, PH3, PCl5,
Examples of stable InCl3, TlBr GeCl4, SnCl2, PbCl2 AsH3, SbH3, SbCl3, BiCl3
compounds

Problem 9.4 : Why Tl⊕1 ion is more stable than Tl⊕3 ?
Solution : Tl is a heavy element which belongs to group 13 of the p-block. The common
oxidation state for this group is 3. In p-block, the lower oxidation state is more stable for
heavier elements due to inert pair effect . Therefore, Tl⊕1 ion is more stable than Tl⊕3 ion.

125

9.4.2 Bonding in compounds of group 13, Table 9.6 : Nature of stable oxides of groups 13,

14 and 15 elements : The lighter elements in 14 and 15 elements

groups 13, 14 and 15 have small atomic radii Group Element Oxide Nature
13
and high ionization enthalpy values. They 14 B B2O3 Acidic
15 Al Al2O3 Amphoteric
form covalent bonds with other atoms by Ga Ga2O3 Amphoteric
In Basic
overlapping of valence shell orbitals. As we Tl In2O3 Basic
C Tl2O3 Acidic
move down the group, the ionization enthalpies
Si CO2 Acidic
are lowered. The atomic radii increase since Ge Acidic
Sn SiO2 Amphoteric
the valence shell orbitals are more diffused. Pb GeO2 Amphoteric
SnO2
The heavier elements in these group tend to N PbO2 Acidic
P Acidic
form ionic bonds. The first member of these As N2O5 Amphoteric
Sb P2O5 Amphoteric
groups belongs to second period and do not Bi As4O6 Basic
Sb2O3
have d orbitals. B, C and N cannot expand Bi2O3

their octet. The subsequent elements in the

group possess vacant d orbital in their valence

shell, which can expand their octet forming a

variety of compounds.

9.4.3 Reactivity towards air/oxygen

a. Group 13 elements : Elements of group 13

on heating with air or oxygen produce oxide of

type E2O3 (where E = element) Do you know ?
4 E (s) + 3O2(g) ∆ 2 E2O3 (s)
2 E (s) + N2 (g) ∆ 2 EN (s) Boron nitride is also called inorganic
b. Group 14 elements : The elements of group graphite.

14 on heating in air or oxygen form oxide 9.4.4 Reaction with water
Most of the elements of groups 13,14 and 15
of the type EO and EO2 in accordance with are unaffected by water. Aluminium reacts
the stable oxidation state and availability of with water on heating and forms hydroxide
while tin reacts with steam to form oxide.
oxygen. 2 Al (s) + 6 H2O (l) ∆ 2 Al (OH)3 (s) + 3 H2 (g)
Sn (s) + 2 H2O (g) ∆ SnO2 (s) + 2 H2 (g)
E (s) + 1/2 O2(g) ∆ EO Lead is unaffected by water, due to formation
E (s) + O2 (g) ∆ EO2 of a protective film of oxide.
c. Group 15 elements : The elements of group

15 on heating in air or oxygen forms two types

of oxide E2O3 and E2O5.
P4 + 3 O2 P4O6
P4 + 5 O2 P4O10
As4 + 3 O2 As4O6 Do you know ?

2Bi + 3 O2 Bi2O3 Phosphorous is stored under water
Increase in metallic character down all these because it catches fire when exposed to air.

groups 13, 14 and 15 reflects their oxides 9.4.4 Reaction with halogens

which gradually vary from acidic through All the elements of group 13 react

amphoteric to basic. (see Table 9.6) directly with halogens to form trihalides (EX3).
Thallium is an exception which forms mono
The nature of stable oxides from groups 13, 14

and 15 are given in Table 9.6. halides (TlX)

2 E (s) + 3 X2 (g) 2 EX3 (s)
All the elements of group 14 (except

carbon) react directly with halogens to form

126

tetrahalides (EX4). The heavy elements Ge The order of catenation of group 14
and Pb form dihalides as well. Stability of di
halides increases down the group. (Refer to elements is C >> Si > Ge = Sn. Lead does not
9.3.1, inert pair effect). The ionic character of
halides also increases steadily down the group. show catenation.This can be clearly seen from

Problem 9.5 : GeCl4 is more stable than the bond enthalpy values.
GeCl2 while PbCl2 is more stable than
PbCl4. Explain. Bond Bond enthalpy (kJmol-1)
Solution : Ge and Pb are the 4th and 5th
period elements down the group 14. The C -C 348
group oxidation state of group 14 is 4 and
the stability of other oxidation state, lower Si - Si 297
by 2 units, increases down the group due
to inert pair effect. The stability of the Ge - Ge 260
oxidation state 2 is more in Pb than in Ge.
Sn - Sn 240
Elements of the group 15 reacts with
halogens to form two series of halides: EX3 and Can you recall?
EX5 . The pentahalides possess more covalent
character due to availability of vacant d orbitals What is common between diamond
of the valence shell for bonding. (Nitrogen and graphite?
being second period element, does not have d
orbitals in its valence shell, and therefore, does 9.6 Allotropy : When a solid element exists
not form pentahalides). Trihalides of the group in different crystalline forms with different
15 elements are predominantly covalent except physical properties such as colour, density,
BiF3. The only stable halide of nitrogen is NF3. melting point, etc. the phenomenon is called
allotropy and individual crystalline forms are
Problem 9.6 : Nitrogen does not form NCl5 called allotropes. Diamond and graphite are
or NF5 but phosphorous can. Explain. well known allotropes of carbon. Fullerenes,
Solution : Phosphorous and other members graphene and carbon nanotubes are other
of the group can make use of d-orbitals in allotropes of carbon. Elements such as boron,
their bonding and thus compounds MX3, as bismuth, silicon, etc. of group 13, 14, and 15
well MX5 are formed Nitrogen can not form exhibit allotropy. In this chapter we are going
NCl5 or NF5 since it is void of d-orbitals in to look at some aspects of allotropes of carbon
its second shell. and phosphorous .
9.6.1 Allotropes of Carbon :
a. Diamond : In diamond each carbon atom
is linked to four other carbon atoms (via. sp3
hybrid orbitals) in tetrahedral manner.

9.5 Catenation : The property of self linking Fig 9.1 Structure of diamond
of atoms of an element by covalent bonds to
form chains and rings is called catenation. The
catenation tendency of an element depends
upon the strength of the bond formed.

Among the group 14 elements,
carbon shows the maximum tendency for
catenation because of the stronger C - C bond
(348 kJ mol-1)

127

The C-C bond distance is 154 pm. when an electric arc is struck between the
The tetrahedra are linked together forming graphite electrodes in an inert atmosphere of
a three dimensional network structure argon or helium. The soot formed contains
(Fig. 9.1) involving strong C-C single bonds. significant amount of C60 fullerene and smaller
Thus diamond is the hardest natural substance amounts of other fullerenes C32, C50, C70 and
with abnormally high melting point (3930 0C). C84.
Diamond is bad conductor of electricity.
Uses : Diamond is used C60 has a shape like soccer ball and called
• for cutting glass and in drilling tools. Buckminsterfullerene or bucky ball. (Fig.
• for making dies for drawing thin wire from 9.3) It contains twenty hexagonal and twelve
pentagonal fused rings of carbons.
metal.
• for making jewellery. Fig. 9.3 Buckminsterfullerene
b. Graphite : Graphite is composed of layers Unlike diamond and graphite, the C60
of two dimensional sheets of carbon atoms fullerene structure exhibits two distinct
(Fig. 9.2) each being made up of hexagonal distances between the neighbouring carbons,
net of sp2 carbons bonded to three neighbours 143.5 pm and 138.3 pm. Fullerenes are
forming three sigma bonds. The fourth electron covalent molecules and soluble in organic
is in the unhybrid p-orbital of each carbon. solvents. Fullerene C60 reacts with group 1
The p-orbitals on all the carbons are parallel metals forming solids such as K3C60. The
to each other. These overlap laterally to form π compound K3C60 behaves as a superconductor
bonds. The π electrons are delocalised over the below 18 K, which means that it conducts
whole layer. The C- C bond length in graphite electric current with zero resistance.
is 141.5 pm. The individual layers are held by d. Cabon nanotubes : Carbon nanotubes are
weak van der Waals forces and separated by cylindrical in shape, consisting of rolled-up
335 pm. This makes graphite soft and slippery. graphite sheet (Fig. 9.4). Nanotubes can be
single-walled (SWNTs) with a diameter of
335 pm less than 1 nm or multi-walled (MWNTs) with

141.5 pm

Fig 9.2 Structure of graphite Carbon atom
Covalent bond
Remember
Fig. 9.4 Carbon nanotubes
Graphite is thermodynamically
most stable form (allotrope) of carbon.

c. Fullerene : Fullerenes are alltropes of
carbon in which carbon molecules are formed
by linking a definite numbers of carbon atoms.
For example : C60. Fullerenes are produced,

128

diameter reaching more than 100 nm. Their • It is translucent white waxy solid.
lengths range from several micrometres to • It glows in dark (chemiluminescence).
millimetres. • It is insoluble in water but dissolves in

Carbon nanotubes are robust. They can be boiling NaOH solution.
bent, and when released, they will spring back • It is poisonous.
to the original shape. b. Red phosphorus : Red phosphorus consists
of chains of P4 tetrahedra linked together by
Carbon nanotubes have high electrical covalent bonds. Thus it is polymeric in nature.
and heat conductivities and highest strength
to weight ratio for any known material to P PP
date. The researchers of NASA are working
on combining carbon nanotubes with other P PP PP P
matertials to obtain composites those can be
used to build light weight space craft. P PP
e. Graphene : Isolated layer of graphite is
called graphene (Fig. 9.5). Graphene sheet is • It is stable and less reactive.
a two dimensional solid. Graphene has unique • It is odourless. It possesses iron grey lustre
electronic properties. • It does not glow in dark.
• It is insoluble in water.
• It is non poisonous.

Remember

Fig. 9.5 Graphene Red phosphorus is prepared by
heating white phosphorus at 573 K in an
Do you know ? inert atmosphere.

The discovery of graphene was 9.7 Molecular structures of some important
awarded with the Nobel prize to Geim and compounds of the group 13, 14 and 15
Novoselov (2010). elements

Can you recall?

9.6.2 Allotropes of phosphorus : Phosphorus • Which element from the following
is found in different allotropic forms, the
important ones being white and red phosphorus. pairs has higher ionization enthalpy?
a. White (yellow) phosphorus : White (yellow) B and Tl, N and Bi
phosphorus consists of discrete tetrahedral P4
molecules. The P-P-P bond angle is 60ο. • Does Boron form covalent compounds

or ionic ?

You have seen in section 9.2 that lighter
P elements of groups 13, 14 and 15 have higher

600 ionization enthalpy and because of smaller
atomic radius do not form cation readily. These

PP elements form covalent compounds. Covalent
molecules have definite shape described with

P the help of bond lengths and bond angles.

White phosphorus is less stable and hence Inorganic molecules are often represented by
more reactive, because of angular strain in the molecular formulae indicating their elemental
composition.
P4 molecule.
129

In the case of covalent inorganic 9.7.4 Diborane (B2H6) : Boron has only three
molecules, the reactivity is better understood valence electron. In diborane each boron atom
from their structures. In this section we will is sp3 hybridized. Three of such hybrid orbitals
consider molecular structures of common are half filled, the fourth sp3 hybrid orbital is
compounds of elements of groups 13,14 and vacant.
15.
9.7.1 Boron trichloride (BCl3) : Boron The two half filled sp3 hybrid orbitals of
trichloride (BCl3) is covalent. Here boron each B atom overlap with 1s orbitals of two
atom is sp2 hybridised having one vacant H atoms and form four B-H covalent bonds.
unhybridised p orbital. B in BCl3 has incomplete Hydrogen atoms are located at the terminal.
octet. The BCl3 is nonpolar trigonal plannar Besides, there are 2-centre - 2-electron bonds
molecule as shown. where ‘1s’ orbital of each of the remaining two
H atoms simultaneouly overlap with half filled
1200 hybrid orbital of one B atom and the vacant
hybrid orbital of the other B atom. This type
9.7.2 Aluminium Chloride (AlCl3) : of overlap produces two three centred - two
Aluminium atom in aluminium chloride electron bonds (3 c - 2e) or banana bonds.
Hydrogen atoms involved in (3 c - 2 e) bonds
is sp2 hybridised, with one vacant unhybrid p are the bridging H- atoms. In diborane two B
orbital. Aluminium Chloride (AlCl3) exists as atoms and four terminal H atoms lie in one
the dimer (Al2Cl6) formed by overlap of vacant plane, while the two bridging H atoms lie
3d orbital of Al with a lone pair of electrons of symmetrically above and below this plane.
Cl.

(3 centred - 2 electron bonds)

9.7.3 Orthoboric acid / boric acid (H3BO3) :
The orthoboric acid has central boron

atom bound to three –OH groups. The solid

orthoboric acid has layered crystal structure in Bonding and structure of diborane
which trigonal planar B(OH)3 units are joined 9.7.5 Silicon dioxide (SiO2) : Silicon dioxide is
together by hydrogen bonds. commonly known as silica. Quartz, cristobalite

and tridymite are different crystalline forms of

silica. They are inter-convertible at a suitable

temperature.

Silicon dioxide (silica) is a covalent

three dimensional network solid, in which

each silicon atom is covalently bound in

tetrahedral manner to four oxygen atoms. The

crystal contains eight membered rings having

Orthoboric acid alternate silicon and oxygen atoms.

130

Borax naturally occurs as tincal (which
contains about 90% borax) in certain inland
lakes in India, Tibet and California (U.S.A).
Preparation :

Borax is prepared from the mineral
colemanite by boiling it with a solution of
sodium carbonate.

Ca2B6O11 + 2 Na2CO3 ∆ Na2B4O7 + 2 NaBO2

Structure of SiO2 Colemanite Borax
9.7.6 Nitric acid (HNO3) : Nitric acid is a
strong, oxidising mineral acid. The central + 2 CaCO3
Properties
nitrogen atom is sp2 hybridised. HNO3 exhibits
resonance phenomenon. i. Borax is white crystalline solid.

H H ii. Borax dissolves in water and gives
OO alkaline solution due to its hydrolysis.
⊕ OO
⊕ Na2B4O7 + 7 H2O 2NaOH + 4 H3BO3
ON
ON

Ortho boric acid

H96pm 1200 Oδ iii. On heating borax first loses water

⊕ 1300 molecules and swells. On further heating
O N140.6pm
Oδ it turns into a transparant liquid, which

solidifies into glass like material known

Resonance Hybrid of HNO3 as borax bead.

9.6.7 Orthophosphoric acid/phosphoric N a2B4O 7.10H2O -1∆0H2O Na2B4O2 7Na∆BO2 + B2O3
acid (H3PO4) : Phosphorus forms number of Borax bead
oxyacids. Orthophosphoric acid is a strong non
toxic mineral acid. It contains three ionizable The borax bead consisting sodium
acidic hydrogens. The central phosphorous
atom is tetrahedral. metaborate and boric anhydride is

used to detect coloured transition metal

ions, under the name borax bead test.

For example, when borax is heated in

a Bunsen burner flame with CoO on a

loop of platinum wire, a blue coloured

Co(BO2)2 bead is formed.
iv. Borax when heated with ethyl alcohol and

Try this concentrated H2SO4 acid, give volatile
vapours of triethyl borate which burn
Find out the structural formulae of
various oxyacids of phosphorus. with green edged flame.

9.8 Chemistry of notable compounds of Na2B4O7 + H2SO4 + 5 H2O Na2SO4 +
elements of groups 13, 14 and 15 4 H3BO3
9.8.1 Borax (Na2B4O7) : It is one of the most H3BO3 + 3C2H5OH B(OC2H5)3 + 3 H2O
important boron compound. The crystalline Triethyl borate
borax has formula Na2B4O7. 10H2O or
Na2[B4O5(OH)4].8H2O. Remember

The above reaction is used as a test
for detection and removal of borate in
qualitative analysis.

131

Uses : Borax is used The chain length of polymer can be controlled
i. in the manufacture of optical and hard by adding (CH3)3 SiCl at the end as shown :

borosilicate glass. ( (CH3 CH3 CH3
ii. as a flux for soldering and welding.
iii. as a mild antiseptic in the preparation of CH3 Si − O − Si − O − Si − CH3
CH3 CH3 n CH3
medical soaps
iv. in quanlitative analysis for borax bead Properties of silicones

test. i. They are water repellant.
v. as a brightner in washing powder.
ii. They have high thermal stability.

9.8.2 Silicones : Silicones represent organo- iii. They are good electrical insulators.

silicon polymers where {R2SiO} is a iv. They are resistant to oxidation and
repeating unit. These are held together by
chemicals.

Uses of silicones

Si O Si linkage. Silicones have They are used as

empirical formula R2SiO (R = CH3 or i. insulating material for electrical appliences

ii. water proofing of fabrics

C6H5 group). This is similar to that of ketone iii. sealant
(R2CO) and hence the name silicones.
iv. high temperature lubricants

Preparation : The starting materials for v. For mixing in paints and enamels to make

manufacture of silicones are alkyl or aryl them resistant to high temperature, sunlight

substituted silicon chlorides, RnSiCl(4 - n) where and chemicals.
R is alkyl or aryl group.
9.8.3 Ammonia (NH3)
When methyl chloride reacts with Preparation

silicon in the presence of copper catalyst at a 1. Ammonia is present in small quantities

temperature 573 K, various types of methyl in air and soil where it is formed by the

substituted chlorosilane of formulae MeSiCl3, decomposition of nitrogeneous organic
Me2SiCl2, Me3SiCl with small amounts of
Me4Si are formed. matter such as urea.

CH3 NH2CONH2 + 2 H2O (NH4)2CO3
Cl − Si − Cl (urea) 2 NH3 + H2O + CO2

2CH3Cl + Si Cu powder 2. Ammonia is prepared on laboratory scale,
573 K
by decomposition of the ammonium salts
CH3
with calcium hydroxide or caustic soda.
CH3 CH3
Cl − Si − Cl 2H2O HO − Si − OH 2NH4Cl + Ca(OH)2 2NH3 + CaCl2 +
-2HCl 2H2O

CH3 CH3 (NH4)2SO4 + 2NaOH 2NH3 + Na2SO4 +
2H2O

CH3 ( (CH3 3. On the large scale ammonia is prepared
n HO − Si − OH
Polymerisation − Si − O − by direct combination of dinitrogen and
CH3 -H2O CH3 n
dihydrogen. (Haber’s process)

Hydrolysis of dimethyl dichlorosilane, N2(g) + 2H2(g) 2NH3(g);∆fHο = -46.1 kJ mol-1

(CH3)2SiCl2 followed by condensation
polymerisation yields straight chain silicone

polymers.

132

High pressure favours the formation of Cu2⊕(aq) + 4NH3(aq) [Cu(NH3)4]2⊕(aq)
ammonia. The optimuim conditions for the
production of ammonia are high pressure deep blue
of 200 × 105 Pa (200 atm), temperature of
∼700 K and use of a catalyst such as iron oxide AgCl(s) + 2NH3(aq) [Ag(NH3)2]Cl(aq)
with trace amounts of K2O and Al2O3. Under colourless
these conditions equilibrium attains rapidly.
Properties Remember
a. Physical properties
i. Ammonia is colourless gas with pungent This reaction is used for the detection
odour. of metal ions such as Cu2+, Ag+.

ii. It has freezing point of 198.4 K and boiling Uses
point of 239.7 K. Ammonia is used in
i. manufacture of fertilizers such as urea,
iii. It is highly soluble in water. The diammonium phosphate, ammonium nitrate,
concentrated aqueous solution of NH3 is ammonium sulphate etc.
called liquor ammonia. ii. manufacture of some inorganic compounds
like nitric acid.
Remember iii. refrigerant (liq. ammonia).
iv. laboratory reagent in qualitative and
Ammonia has higher melting point quantitative analysis (aq. solution of ammonia)
and boiling point, because in the solid and
liquid state NH3 molecules get associated Remember
together through hydrogen bonding. Thus
some more energy is required to break such Ammonia gives brown ppt with
intermolecular hydrogen bonds.
Nessler’s reagent (alkaline solution of

b. Chemical properties K2HgI4).

i. Basic nature : The aqueous solution of 2 KI + HgCl2 HgI2 + 2 KCl

ammonia is basic in nature due to the formation 2 KI + HgI2 K2HgI4 (Nesseler’s reagent)

of OH− ions. 2 K2HgI4 + NH3 + 3 KOH H2N-HgO- HgI

NH3 (g) + H2O (l) NH⊕4 (aq) + OH (aq) Millon’s base (Brown ppt)
ii. Ammonia reacts with acids to form
+7 KI + 2 H2O
ammonium salts.

NH3 + HCl NH4Cl
2NH3 + H2SO4 (NH4)2SO4
iii. Aqueous solution of ammonia precipitates

out as hydroxides (or hydrated oxides) of

metals from their salt solutions.

ZnSO4(aq) + 2 NH4OH(aq) Zn(OH)2(s) +

White ppt

FeCl3(aq) + 2NH4OH(aq) (NH4)2SO4(aq)
Fe2O3.xH2O +

iv. Complex formation : NH4Cl(aq)

The lone pair of electrons on nitrogen the atom

facilitates complexation of ammonia with

transition metal ions.

133

Exercises

1. Choose correct option. 8. Find out the difference between

A. Which of the following is not an A. Diamond and Graphite

allotrope of carbon ? B. White phosphorus and Red phosphorus

a. bucky ball b. diamond 9. What are silicones ? Where are they used ?

c. graphite d. emerald 10. Explain the trend in oxidation state of

B. is inorganic graphite elements from nitrogen to bismuth.

a. borax b. diborane 11. Give the test that is used to detect borate

c. boron nitride d. colemanite radical is qualitative analysis.

C. Haber’s process is used for preparation 12. Explain structure and bonding of

of diborane.

a. HNO3 b. NH3 13. A compound is prepared from the
c. NH2CONH2 d. NH4OH
D. Thallium shows different oxidation mineral colemanite by boiling it with a

state because solution of sodium carbonate. It is white

a. of inert pair effect crystalline solid and used for inorganic

b. it is inner transition element qualitative analysis.

c. it is metal a. Name the compound produced.

d. of its high electronegativity b. Write the reaction that explains its

E. Which of the following shows most formation.

prominent inert pair effect ? 14. Ammonia is a good complexing agent.

a. C b. Si Explain.

c. Ge d. Pb 15. State true or false. Correct the false

2. Identify the group 14 element that best statement.

fits each of the following description. A. The acidic nature of oxides of group 13

A. Non metallic element increases down the graph.

B. Form the most acidic oxide B. The tendency for cantenation is much

C. They prefer +2 oxidation state. higher for C than for Si.

D. Forms strong π bonds. 16. Match the pairs from column A and B.

3. Give reasons. A B

A. Ga3⊕ salts are better reducing agent BCl3 Angular molecule
SiO2 linear covalent molecule
while Tl3⊕ salts are better oxidising CO2 Tetrahedral molecule
Planar trigonal molecule
agent.

4. BG.ivPebCtlh4 eis less stable than PcobmClp2 ound in 17. Give the reactions supporting basic
formula of a
nature of ammonia.
which carbon exhibit an oxidation state
18. Shravani was performing inorganic
of
qualitative analysis of a salt. To an
A. +4 B. +2 C. -4
aqueous solution of that salt, she added
5. Explain the trend of the following in
silver nitrate. When a white precipitate
group 13 elements :
was formed. On adding ammonium
A. atomic radii B. ionization enthalpy
hydroxide to this, she obtained a clear
C. electron affinity
solution. Comment on her observations
6. Answer the following
and write the chemical reactions
A. What is hybridization of Abal ninanAalCbol3n?d.
B. Name a molecule having involved.

7. Draw the structure of the following Activity :

A. Orthophosphoric acid Prepare models of allotropes of carbon
B. Resonance structure of nitric acid and phosphorous.

134

10. States of Matter : Gaseous and Liquid States

10.1 Introduction:
We have learnt that substances exist

in one of the three main states of matter. The
three distinct physical forms of a substance are
Solid, Liquid, and Gas.

Can you recall? Fig. 10.1 : Different States of Water

Water exists in the three different Three states of matter are interconvertible
forms solid ice, liquid water and gaseous by exchange of Heat as given below:
vapours.
Solid Heat Liquid Heat Gas
Key points of differentiation between Cool
the three states can be understood as given in Cool
Table 10.1.

Table 10.1: Distinguishing points between Solid, Liquid and Gas

Sr. No. Points Solid Liquid Gas
1 Microscopic view

Mean atomic/ Mean separation ≈ Mean separation ≈ Mean separation >
molecular 3-5A0 3-10A0 5A0
separation

2 Arrangement of Particles are tightly Particles are loosely Particles are more

particles (atoms/ held, and have regular packed, irregular loosely packed,

molecules) arrangement of atoms/ arragement of particles highly irregular

molecules arrangement

3 Movement of Particles cannot move Particles can move a Particles are in
particles
freely as they occupy small distance within the continuous

fixed positions. liquid random motion.

4 Shape and Has definite shape and Takes the shape of Takes the shape and
volume the volume of its
volume the container and has container.

definite volume.

5 Intermoleculer Very small Moderate Intermolecular Large

space Intermolecular space space Intermolecular

space.

6 Effect of a Volume change is Moderate effect on Volume change
volume change significantly high.
small change in small

temperature

7 Compression or Practically Small Compressibility Compressible

Expansion Non- compressible

135

10.2 Intermolecular Forces : Intermolecular Dipole moment (µ) is the product of the
forces are the attractive forces as well magnitude of the charge (Q) and the distance
as repulsive forces present between the between the centres of positive and negative
neighbouring molecules. The attractive force charge (r). It is designated by a Greek Letter
decreases with the increase in distance between (µ) (mu). Its unit is debye (D).
the molecules. The intermolecular forces are
strong in solids, less strong in liquids and very Dipole moment is a vector
weak in gases. Thus, the three physical states quantity and is depicted by a small arrow
of matter can be determined as per the strength with tail in the positive centre and head
of intermolecular forces. pointing towards the negative centre.

The physical properties of matter such as δ⊕ µ δ
melting point, boiling point, vapor pressure,
viscosity, evaporation, surface tension and H Cl
solubility can be studied with respect to
the strength of attractive forces between More Charge density
the molecules. During the melting process towards chlorine
intermolecular forces are partially overcome,
whereas they are overcome completely during Unequal sharing of electrons
the vapourization process. between the bonding atoms.
10.2.1 Types of Intermolecular Forces: Fig 10.2 (a) : Polar molecule
The four types of intermolecular forces are-
i. Dipole-dipole interactions δI ⊕ Cδl δI ⊕ Cδl δI⊕
ii. Ion-dipole interactions Cδl δI⊕ Cl δI ⊕ Cδl
iii. Dipole-Induced dipole interaction
iv. London Dispersion Forces δ
v. Hydrogen bonding Fig 10.2 (b) : Dipole-dipole interaction
i. Dipole-dipole interactions : Polar
molecules experience dipole-dipole forces due For example, the dipole moment of HF
to electrostatic interactions between dipoles on may be represented as: H-F
neighboring molecules.
What are polar molecules or Polar covalent The crossed arrow( ) above the Lewis
molecules? (Refer to Chapter 5). structure represents an electron density shift.

Polar covalent molecule is also described Thus polar molecules have permanent
as “dipole” meaning that the molecule has two dipole moments. When a polar molecule
‘poles’. The covalent bond becomes polar due encounters another polar molecule, the positive
to electronegativity difference between the end of one molecule is attracted to the negative
bonding atoms. Hence polarity is observed in end of another polar molecule. Many such
the compounds containing dissimilar atoms. molecules have dipoles and their interaction
For example, HCl molecule (see Fig. 10.2 (a)). is termed as dipole-dipole interaction. These
forces are generally weak, with energies of the
One end (pole) of the molecule has order of 3-4 kJ mol-1 and are significant only
partial positive charge on hydrogen atom when molecules are in close contact. i.e. in a
while at other end chlorine atom has partial solid or a liquid state.
negative charge (denoted by Greek letter ‘δ’
delta). As a result of polarisation, the molecule For example C4H9Cl, (butyl chloride),
possesses the dipole moment. CH3 - O - CH3 (dimethyl ether) ICl (iodine
chloride B.P. 27 0C), are dipolar liquids.
The molecular orientations due to dipole-
dipole interaction in ICl liquid is shown in
Fig. 10.2 (b).

136

In brief, more polar the substance, greater Cations are smaller in size than the
isoelectronic anions. The charge density on
the strength of its dipole-dipole interactions. cation (Na⊕) is more concentrated than anion
(Cl ). This makes the interaction between (Na⊕)
Table 10.2 enlists several substances with and negative end of the polar H2O molecule
(Fig. 10.2 (c)) stronger than the corresponding
similar molecular masses but different dipole interaction between (Cl ) and positive end of
the polar H2O molecule.
moments. From Table 10.2, it is clear that
More the charge on cation, stronger is
higher the dipole moment, stronger are the the ion-dipole interaction. For example, Mg2⊕
ion has higher charge and smaller ionic radius
inter molecular forces, generally leading to (78 pm) than Na⊕ ion (98 pm), hence Mg2⊕ ion
is surrounded (hydrated) more strongly with
higher boiling points. water molecules and exerts strong ion-dipole
interaction.
Table 10.2 : Effect of dipole moments on boiling Thus the strength of interaction
increases with increase in charge on cation and
point (b.p.) with decrease in ionic size or radius. Therefore,
ion-dipole forces increase in the order :
Substance Molar Dipole b.p. Na⊕ < Mg2⊕ < Al3⊕.
Mass Moment (K) iii. Dipole-Induced dipole interaction :
(amu) (D)
When polar molecules (like H2O,
CH3 - CH2 - CH3 44.10 0.1 231 NH3) and nonpolar molecules (like benzene)
approach each other, the polar molecules
CH3 - O - CH3 46.07 1.3 248 induce dipole in the non-polar molecules.
Hence ‘Temporary dipoles’ are formed
CH3 - Cl 50.49 1.9 249 by shifting of electron clouds in nonpolar
molecules. For example, Ammonia (NH3)
CH3 - CN 41.05 3.9 355 is polar and has permanent dipole moment
while Benzene (C6H6) is non polar and has
When different substances coexist zero dipole moment. The force of attraction
developed between the polar and nonpolar
in single phase, following intermolecular molecules is of the type dipole - induced dipole
interaction. It can be seen in Fig 10.2(d) in the
interactions are present. following manner:

ii. Ion-dipole interactions : An ion-dipole

force is the result of electrostatic interactions

between an ion (cation or anion) and the

partial charges on a polar molecule.

The strength of this interaction depends

on the charge and size of an ion. It also depends

on the magnitude of dipole moment and size of

the molecule.

(Hydrated Na⊕ ion)

o
oo

oo
o

Fig. 10.2 (c) : Na⊕ ion(cation) - H2O interaction Polar Non - Polar Non - Polar molecular
Ion-dipole forces are particularly molecular molecular with induced dipole

important in aqueous solutions of ionic Fig. 10.2 (d) : Dipole - induced dipole
substances such as sodium chloride (NaCl). interaction
When an ionic compound, sodium chloride is
dissolved in water, the ions get separated and
surrounded by water molecules which is called
Hydration of sodium ions.

137

iv. London Dispersion Force : The study molecule, higher is its ability to become

of intermolecular forces present among polar. Similarly more the spread out shapes,

nonpolar molecules or the individual atoms of higher the dispersion forces present between

a noble gas is very interesting. For example, the molecules. London dispersion forces

Benzene (C6H6) has zero dipole moment and are stronger in a long chain of atoms where
experiences no dipole-dipole forces, yet exists molecules are not compact. This can affect

in liquid stage. physical property such as B.P. n-Pentane,

In case of nonpolar molecules and inert for example boils at 309.4 K, whereas

gases, only dipersion forces exist. Dispersion neo - pentane boils at 282.7 K. Here both the

forces are also called as London forces due substances have the same molecular formula,
to an idea of momentary dipole which was
proposed by the German Physicist, Fritz C5H12, but n-pentane is longer and somewhat
London in 1930. These forces are also called spread out, where as neo-pentane is more

as van der Waals forces. It is the weakest Fig 10.3 (a):
intermolecular force that develops due to Longer, less
interaction between two nonpolar molecules. compact molecule

In general, all atoms and molecules experience H H H H H n-Pentane
London dispersion forces, which result from
the motion of electrons. At any given instant of H C C C C C H (b.p.= 309.4K)
time, the electron distribution in an atom may
H H H H H H

be asymmetrical, giving the atom a short lived H C H
dipole moment. This momentary dipole on one HH

atom can affect the electron distribution in the HC C CH

neighbouring atoms and induce momentary HH
dipoles in them. As a result, weak attractive H C H
force develops.
H
For example, substances composed of Fig 10.3 (b): More compact molecule
molecules such as O2, CO2, N2, halogens,
methane gas, helium and other noble gases neo - pentane (b.p.= 282.7K)
show van der Waals force of attraction.
spherical and compact (see Fig. 10.3).
The strength of London forces increases v. Hydrogen Bonding : A hydrogen bond is a
with increase in molecular size, molecular special type of dipole-dipole attraction which
mass and number of electrons present in an occurs when a hydrogen atom is bonded to a
atom or molecule. strongly electronegative atom or an atom with
a lone pair of electrons. Hydrogen bonds are
When two nonpolar molecules approach generally stronger than usual dipole - dipole
each other, attractive and repulsive forces and dispersion forces, and weaker than true
between their electrons and the nuclei will lead covalent or ionic bonds.
to distortions in the size of electron cloud, a Definition : The electrostatic force of attraction
property referred to as polarizability. between positively polarised hydrogen atom
of one molecule and a highly electronegative
Polarizability is a measure of how easily atom (which may be negatively charged) of
an electron cloud of an atom is distorted by
an applied electric field. It is the property of other molecule is called as hydrogen bond.

atom. The ability to form momentary dipoles Strong electronegative atoms that

that means the ability of another molecule to form hydrogen bonds are nitrogen, oxygen,

become polar by redistributing its electrons and fluorine. Hydrogen bond is denoted

is known as polarizability of the atom or by ( ) dotted line. Hydrogen bond which

molecule. occurs within the same molecule represents

More the number of electrons in a Intramolecular Hydrogen bond.

138

A hydrogen bond present between electron pairs on oxygen atom.
two like or unlike molecules, represents The boiling point generally increases
Intermolecular Hydrogen bond.
(See Fig. 10.4 (a) and (b)). with increase in molecular mass, but the
i. Inter molecular H-bonding hydrides of nitrogen (NH3), oxygen (H2O) and
H-bonding in H−F : fluorine (HF) have abnormally high boiling
Hδ+−Fδ- Hδ+−Fδ- Hδ+−Fδ- Hδ+−Fδ- points due to the presence of hydrogen bonding
between the molecules.
H-bond
Similarly due to presence of H- bond,
Fig. 10.4 : H-bonding in H2O (above) and viscosity of liquid increases. Hydrogen bonds
NH3 (below) (solid line denotes covalent bond, play vital role in determining structure and
dotted line denotes hydrogen bond) properties of proteins and nucleic acid present
ii. Intra molecular H-bonding : in all living organisms. In case of gases
intermolecular forces of attraction are very
H-bonding in ethylene glycol : weak.
CH2−Oδ-−Hδ+ 10.2.2 Intermolecular Forces and Thermal
CH2−Oδ-−Hδ+ energy : Thermal energy is the origin of
H - bonding and boiling point : Due to the kinetic energy of the particles of matter that
presence of H-bonding in the compounds, arises due to movement of particles (about
more energy is required to break the bonds. which you must have learnt in Physics). It is
Therefore, boiling point is more in case of directly proportional to the temperature; that
liquid molecules containing H-bond. Hydrogen means thermal energy increases with increase
bonds can be quite strong with energies up to in temperature and vice versa.
40 kJ/mol.
Water in particular is able to form a vast Three states of matter are the
three dimensional network of hydrogen bonds consequence of a balance between the
as shown in Fig. 10.5. It is because each H2O intermolecular forces of attraction and the
molecule has two hydrogen atoms and two thermal energy of the molecules.

If the intermolecular forces are very
weak, molecules do not come together to
make liquid or solid unless thermal energy is
decreased by lowering the temperature.
When a substance is to be converted
from its gaseous state to solid state, its thermal
energy (or temperature) has to be reduced. At
this stage, the intermolecular forces become
more important than thermal energy of the
particles.

Intermolecular Force Increases

Gas Liquid Solid

Thermal energy decreases

Fig. 10.5 : Liquid water contains a vast A comparison of various kinds of
three dimensional network of hydrogen intermolecular forces discussed in this section
is made in Table 10.3.
bonds

139

Table 10.3 : Comparison of Intermolecular Forces

Force Strength Characterstics
Ion-dipole
Dipole-dipole Moderate (10 - 50 kJ/mol) Occurs between ions and polar solvents
London dispersion
Weak (3 - 4 kJ/mol) Occurs between polar molecules
Hydrogen bond
Weak (1 - 10 kJ/mol) Occurs between all molecules; strength
depends on size, polarizability

Moderate (10 - 40 kJ/mol) Occurs between molecules with O-H, N-H,
and H-F bonds

Can you tell? Do you know ?
What are the various components
present in the atmosphere? Among the three compounds H2, H2S
Name five elements and five and H2Se, the first one, H2O has the smallest
compounds those exist as gases at room molecular mass. But it has the highest B.P.
temperature. of 1000C. B.P. of H2S is -600C and of H2Se is
-41.25 0C. The extraordinary high B.P. of H2O
Just think is due to very strong hydrogen bonding even
though it has the lowest molecular mass.
What is air?
It is a mixture of various gases. Air, 10.3 Characteristic properties of Gases :
we can not see but feel the cool breeze. The Under normal conditions, out of 118
composition of air by volume is around 78
percent N2, 21 percent O2 and 1 percent elements from the periodic table, only a few
other gases including CO2. The chemistry (eleven) elements exist as gases. The gaseous
of atmospheric gases is an important state is characterized by the following physical
subject of study as it involves air pollution. properties :
O2 in air is essential for survival of aerobic 1. Gases are lighter than solids and liquids i.e.
life. possess lower density.
2. Gases do not possess a fixed volume and
Nitogen 78% shape. They occupy entire space available and
take the shape of the container.
Oxygen 21% 3. Gas molecules are widely seperated and
Carbon are in continuous, random motion. Therefore,
Dioxide and gases exert pressure equally in all directions
other gases due to collision of gas molecules, on the walls
0.03% of the container.
Inert gases 4. In case of gases, intermolecular forces are
(mainly weakest.
argon) 0.97% 5. Gases possess the property of diffusion
Water vapor which is a spontaneous homogenous inter
mixing of two or more gases.
Composition of air with respect to 6. Gases are highly compressible.
proportion of various gases Measurable properties of Gases : (Refer
to Chapter 1) : Some Important measurable
properties of the gases are given below :
1. Mass: The mass, m, of a gas sample is
measure of the quantity of matter it contains.
It can be measured experimentally. SI unit of
mass of gas is kilogram (kg). 1kg = 103g.

140

The mass of a gas is realated to the 4. Temperature : It is the property of an object

number of moles (n) by the expression i. that determines direction in which energy will

=i. n m=olmarasmsainssgirnagmrsams m flow when that object is in contact with other
M
object.

The following expressions are also useful in In scientific measurements temperature

calculations. (T) is measured either on the celsius scale (0C)

ii. or the Kelvin scale (K). (Note that degree sign

n Number of molecules N N is not used in Kelvin unit).
Avogadro Number NA 6.022 u1023
SI Unit of Temperature of a gas is kelvin

iii. (K). The celsius and kelvin scales are related
n=
Volume of a given gas in litres at STP by the expression
22.414 litres at STP
T (K) = t0 C + 273.15

2. Volume: Volume (V) of a sample of gas is 5. Density : It is the mass per unit volume.

the amount of space it occupies. It is expressed d= m
V
in terms of different units like Litres (L),
∴ SI Unit of density is kg m-3.
mililitres(mL), cubic centimeter (cm3), cubic
In the case of gases, relative density is
meter (m3) or decimeter cube (dm3). SI Unit of
measured with respect to hydrogen gas and is
the volume is cubic meter (m3).
called vapour density.
Most commonly used unit to measure molar mass
∴ Vapour density = 2
the volume of the a gas is decimeter cube or

litre. 6. Diffusion: In case of gases, Rate of diffusion
of two or more gases is measured.
1 L = 1000mL = 1000 cm3 = 1dm3

1 m3 = 103 dm3 = 103L = 106cm3 = 106 mL

3. Pressure: Pressure (P) is defined as force Rate Volume of a gas diffused
of = time required for diffusion
per unit area. Force f diffusion
Area a
Pressure = =

Pressure of gas is measured with ∴ SI Unit for Rate of diffusion is dm3 s-1 or
cm3 s-1
‘manometer’ and atmospheric pressure is 10.4 Gas Laws : The behavior of gases can
be studied by four variables namely pressure,
measured by ‘barometer’. volume,temperature and the number of moles.
These variables and measurable properties of
SI Unit of pressure is pascal (Pa) or the gases are related with one another through
different gas laws.
Newton per meter square (N m-2).
Think of a gas in a cylinder or a sealed
1 Pa = 1 Nm-2 = 1 kg m-1 s-2 container. We can measure number of moles
(n) of gas inside, the pressure (P) of a gas,
1 bar = 1.00 × 105 Pa the volume (V) of a gas (which is equal to the
volume of the container) and the temperature
The bar is now replacing the standard (T) of the gas. The observed relationships
between P, V, n and T are summarized by
atmosphere (atm) as the most convenient unit five gas laws: Boyle’s law, Charles’ law, Gay
Lussac law, Avogadro law and Dalton’s law.
of pressure.

1 atm = 76 cm of Hg = 760 mm of Hg =

760 torr as 1 torr = 1mm Hg

1 atm = 101.325 kPa = 101325 Pa =

1.01325 × 105 N m-2 = 1.01325 bar.

Can you tell?

What is the unit in which car-tyre
pressure is measured ?

141