NCERT Textbook - The s Block Elements Class 11 Notes | EduRev

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Class 11 : NCERT Textbook - The s Block Elements Class 11 Notes | EduRev

 Page 1


291       C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd  Reprint   27.7.6
291 THE s-BLOCK ELEMENTS
The s-block elements of the Periodic Table are those in
which the last electron enters the outermost s-orbital. As
the s-orbital can accommodate only two electrons, two
groups (1 & 2) belong to the s-block of the Periodic Table.
Group 1 of the Periodic Table consists of the elements:
lithium, sodium, potassium, rubidium, caesium and
francium. They are collectively known as the alkali metals.
These are so called because they form hydroxides on
reaction with water which are strongly alkaline in nature.
The elements of Group 2 include beryllium, magnesium,
calcium, strontium, barium and radium. These elements
with the exception of beryllium are commonly known as
the alkaline earth metals. These are so called because their
oxides and hydroxides are alkaline in nature and these
metal oxides are found in the earth’s crust*.
Among the alkali metals sodium and potassium are
abundant and lithium, rubidium and caesium have much
lower abundances (Table 10.1). Francium is highly
radioactive; its longest-lived isotope 
223
Fr has a half-life of
only 21 minutes. Of the alkaline earth metals calcium and
magnesium rank fifth and sixth in abundance respectively
in the earth’s crust. Strontium and barium have much
lower abundances. Beryllium is rare and radium is the
rarest of all comprising only 10
–10
 per cent of igneous
rocks
†
 (Table 10.2, page 299).
The general electronic configuration of s-block elements
is [noble gas]ns
1
 for alkali metals and [noble gas] ns
2
 for
alkaline earth metals.
UNIT 10
After studying this unit, you will be
able to
• • • • • describe the general charact-
eristics of the alkali metals and
their compounds;
• • • • • explain the general characteristics
of the alkaline earth metals and
their compounds;
• • • • • describe the manufacture,
properties and uses of industrially
important sodium and calcium
compounds including Portland
cement;
• • • • • appreciate the biological
significance of sodium,
potassium, magnesium and
calcium.
THE s -BLOCK ELEMENTS
* The thin, rocky outer layer of the Earth is crust.   † A type of rock formed
from magma (molten rock) that has cooled and hardened.
The first element of alkali and alkaline earth metals differs
in many respects from the other members of the group
© NCERT
not to be republished
Page 2


291       C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd  Reprint   27.7.6
291 THE s-BLOCK ELEMENTS
The s-block elements of the Periodic Table are those in
which the last electron enters the outermost s-orbital. As
the s-orbital can accommodate only two electrons, two
groups (1 & 2) belong to the s-block of the Periodic Table.
Group 1 of the Periodic Table consists of the elements:
lithium, sodium, potassium, rubidium, caesium and
francium. They are collectively known as the alkali metals.
These are so called because they form hydroxides on
reaction with water which are strongly alkaline in nature.
The elements of Group 2 include beryllium, magnesium,
calcium, strontium, barium and radium. These elements
with the exception of beryllium are commonly known as
the alkaline earth metals. These are so called because their
oxides and hydroxides are alkaline in nature and these
metal oxides are found in the earth’s crust*.
Among the alkali metals sodium and potassium are
abundant and lithium, rubidium and caesium have much
lower abundances (Table 10.1). Francium is highly
radioactive; its longest-lived isotope 
223
Fr has a half-life of
only 21 minutes. Of the alkaline earth metals calcium and
magnesium rank fifth and sixth in abundance respectively
in the earth’s crust. Strontium and barium have much
lower abundances. Beryllium is rare and radium is the
rarest of all comprising only 10
–10
 per cent of igneous
rocks
†
 (Table 10.2, page 299).
The general electronic configuration of s-block elements
is [noble gas]ns
1
 for alkali metals and [noble gas] ns
2
 for
alkaline earth metals.
UNIT 10
After studying this unit, you will be
able to
• • • • • describe the general charact-
eristics of the alkali metals and
their compounds;
• • • • • explain the general characteristics
of the alkaline earth metals and
their compounds;
• • • • • describe the manufacture,
properties and uses of industrially
important sodium and calcium
compounds including Portland
cement;
• • • • • appreciate the biological
significance of sodium,
potassium, magnesium and
calcium.
THE s -BLOCK ELEMENTS
* The thin, rocky outer layer of the Earth is crust.   † A type of rock formed
from magma (molten rock) that has cooled and hardened.
The first element of alkali and alkaline earth metals differs
in many respects from the other members of the group
© NCERT
not to be republished
292       C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd  Reprint   27.7.6
292 CHEMISTRY
Lithium and beryllium, the first elements
of Group 1 and Group 2 respectively exhibit
some properties which are different from those
of the other members of the respective group.
In these anomalous properties they resemble
the second element of the following group.
Thus, lithium shows similarities to magnesium
and beryllium to aluminium in many of their
properties. This type of diagonal similarity is
commonly referred to as diagonal relationship
in the periodic table. The diagonal relationship
is due to the similarity in ionic sizes and /or
charge/radius ratio of the elements.
Monovalent sodium and potassium ions and
divalent magnesium and calcium ions are
found in large proportions in biological fluids.
These ions perform important biological
functions such as maintenance of ion balance
and nerve impulse conduction.
10.1 GROUP 1 ELEMENTS: ALKALI
METALS
The alkali metals show regular trends in their
physical and chemical properties with the
increasing atomic number. The atomic,
physical and chemical properties of alkali
metals are discussed below.
10.1.1 Electronic Configuration
All the alkali metals have one valence electron,
ns
1
 (Table 10.1) outside the noble gas core.
The loosely held s-electron in the outermost
valence shell of these elements makes them the
most electropositive metals. They readily lose
electron to give monovalent M
+
 ions. Hence they
are never found in free state in nature.
increase in atomic number, the atom becomes
larger. The monovalent ions (M
+
) are smaller
than the parent atom. The atomic and ionic
radii of alkali metals increase on moving down
the group i.e., they increase in size while going
from Li to Cs.
10.1.3 Ionization Enthalpy
The ionization enthalpies of the alkali metals
are considerably low and decrease down the
group from Li to Cs. This is because the effect
of increasing size outweighs the increasing
nuclear charge, and the outermost electron is
very well screened from the nuclear charge.
10.1.4 Hydration Enthalpy
The hydration enthalpies of alkali metal ions
decrease with increase in ionic sizes.
Li
+
> Na
+
 > K
+
 > Rb
+
 > Cs
+
Li
+
 has maximum degree of hydration and
for this reason lithium salts are mostly
hydrated, e.g., LiCl· 2H
2
O
10.1.5 Physical Properties
All the alkali metals are silvery white, soft and
light metals. Because of the large size, these
elements have low density which increases
down the group from Li to Cs. However,
potassium is lighter than sodium. The melting
and boiling points of the alkali metals are low
indicating weak metallic bonding due to the
presence of only a single valence electron in
them. The alkali metals and their salts impart
characteristic colour to an oxidizing flame. This
is because the heat from the flame excites the
outermost orbital electron to a higher energy
level. When the excited electron comes back to
the ground state, there is emission of radiation
in the visible region as given below:
Alkali metals can therefore, be detected by
the respective flame tests and can be
determined by flame photometry or atomic
absorption spectroscopy. These elements when
irradiated with light, the light energy absorbed
may be sufficient to make an atom lose electron.
Element Symbol Electronic configuration
Lithium Li 1s
2
2s
1
Sodium Na 1s
2
2s
2
2p
6
3s
1
Potassium K 1s
2
2s
2
2p
6
3s
2
3p
6
4s
1
Rubidium Rb 1s
2
2s
2
2p
6
3s
2
3p
6
3d
10
4s
2
4p
6
5s
1
Caesium Cs 1s
2
2s
2
2p
6
3s
2
3p
6
3d
10
4s
2
4p
6
4d
10
5s
2
5p
6
6s
1 
or [Xe] 6s
1
Francium Fr [Rn]7s
1
10.1.2 Atomic and Ionic Radii
The alkali metal atoms have the largest sizes
in a particular period of the periodic table. With
Metal Li Na K Rb Cs
Colour Crimson Yellow Violet Red Blue
red violet
?/nm 670.8 589.2 766.5 780.0 455.5
© NCERT
not to be republished
Page 3


291       C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd  Reprint   27.7.6
291 THE s-BLOCK ELEMENTS
The s-block elements of the Periodic Table are those in
which the last electron enters the outermost s-orbital. As
the s-orbital can accommodate only two electrons, two
groups (1 & 2) belong to the s-block of the Periodic Table.
Group 1 of the Periodic Table consists of the elements:
lithium, sodium, potassium, rubidium, caesium and
francium. They are collectively known as the alkali metals.
These are so called because they form hydroxides on
reaction with water which are strongly alkaline in nature.
The elements of Group 2 include beryllium, magnesium,
calcium, strontium, barium and radium. These elements
with the exception of beryllium are commonly known as
the alkaline earth metals. These are so called because their
oxides and hydroxides are alkaline in nature and these
metal oxides are found in the earth’s crust*.
Among the alkali metals sodium and potassium are
abundant and lithium, rubidium and caesium have much
lower abundances (Table 10.1). Francium is highly
radioactive; its longest-lived isotope 
223
Fr has a half-life of
only 21 minutes. Of the alkaline earth metals calcium and
magnesium rank fifth and sixth in abundance respectively
in the earth’s crust. Strontium and barium have much
lower abundances. Beryllium is rare and radium is the
rarest of all comprising only 10
–10
 per cent of igneous
rocks
†
 (Table 10.2, page 299).
The general electronic configuration of s-block elements
is [noble gas]ns
1
 for alkali metals and [noble gas] ns
2
 for
alkaline earth metals.
UNIT 10
After studying this unit, you will be
able to
• • • • • describe the general charact-
eristics of the alkali metals and
their compounds;
• • • • • explain the general characteristics
of the alkaline earth metals and
their compounds;
• • • • • describe the manufacture,
properties and uses of industrially
important sodium and calcium
compounds including Portland
cement;
• • • • • appreciate the biological
significance of sodium,
potassium, magnesium and
calcium.
THE s -BLOCK ELEMENTS
* The thin, rocky outer layer of the Earth is crust.   † A type of rock formed
from magma (molten rock) that has cooled and hardened.
The first element of alkali and alkaline earth metals differs
in many respects from the other members of the group
© NCERT
not to be republished
292       C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd  Reprint   27.7.6
292 CHEMISTRY
Lithium and beryllium, the first elements
of Group 1 and Group 2 respectively exhibit
some properties which are different from those
of the other members of the respective group.
In these anomalous properties they resemble
the second element of the following group.
Thus, lithium shows similarities to magnesium
and beryllium to aluminium in many of their
properties. This type of diagonal similarity is
commonly referred to as diagonal relationship
in the periodic table. The diagonal relationship
is due to the similarity in ionic sizes and /or
charge/radius ratio of the elements.
Monovalent sodium and potassium ions and
divalent magnesium and calcium ions are
found in large proportions in biological fluids.
These ions perform important biological
functions such as maintenance of ion balance
and nerve impulse conduction.
10.1 GROUP 1 ELEMENTS: ALKALI
METALS
The alkali metals show regular trends in their
physical and chemical properties with the
increasing atomic number. The atomic,
physical and chemical properties of alkali
metals are discussed below.
10.1.1 Electronic Configuration
All the alkali metals have one valence electron,
ns
1
 (Table 10.1) outside the noble gas core.
The loosely held s-electron in the outermost
valence shell of these elements makes them the
most electropositive metals. They readily lose
electron to give monovalent M
+
 ions. Hence they
are never found in free state in nature.
increase in atomic number, the atom becomes
larger. The monovalent ions (M
+
) are smaller
than the parent atom. The atomic and ionic
radii of alkali metals increase on moving down
the group i.e., they increase in size while going
from Li to Cs.
10.1.3 Ionization Enthalpy
The ionization enthalpies of the alkali metals
are considerably low and decrease down the
group from Li to Cs. This is because the effect
of increasing size outweighs the increasing
nuclear charge, and the outermost electron is
very well screened from the nuclear charge.
10.1.4 Hydration Enthalpy
The hydration enthalpies of alkali metal ions
decrease with increase in ionic sizes.
Li
+
> Na
+
 > K
+
 > Rb
+
 > Cs
+
Li
+
 has maximum degree of hydration and
for this reason lithium salts are mostly
hydrated, e.g., LiCl· 2H
2
O
10.1.5 Physical Properties
All the alkali metals are silvery white, soft and
light metals. Because of the large size, these
elements have low density which increases
down the group from Li to Cs. However,
potassium is lighter than sodium. The melting
and boiling points of the alkali metals are low
indicating weak metallic bonding due to the
presence of only a single valence electron in
them. The alkali metals and their salts impart
characteristic colour to an oxidizing flame. This
is because the heat from the flame excites the
outermost orbital electron to a higher energy
level. When the excited electron comes back to
the ground state, there is emission of radiation
in the visible region as given below:
Alkali metals can therefore, be detected by
the respective flame tests and can be
determined by flame photometry or atomic
absorption spectroscopy. These elements when
irradiated with light, the light energy absorbed
may be sufficient to make an atom lose electron.
Element Symbol Electronic configuration
Lithium Li 1s
2
2s
1
Sodium Na 1s
2
2s
2
2p
6
3s
1
Potassium K 1s
2
2s
2
2p
6
3s
2
3p
6
4s
1
Rubidium Rb 1s
2
2s
2
2p
6
3s
2
3p
6
3d
10
4s
2
4p
6
5s
1
Caesium Cs 1s
2
2s
2
2p
6
3s
2
3p
6
3d
10
4s
2
4p
6
4d
10
5s
2
5p
6
6s
1 
or [Xe] 6s
1
Francium Fr [Rn]7s
1
10.1.2 Atomic and Ionic Radii
The alkali metal atoms have the largest sizes
in a particular period of the periodic table. With
Metal Li Na K Rb Cs
Colour Crimson Yellow Violet Red Blue
red violet
?/nm 670.8 589.2 766.5 780.0 455.5
© NCERT
not to be republished
293       C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd  Reprint   27.7.6
293 THE s-BLOCK ELEMENTS
Property Lithium Sodium Potassium Rubidium Caesium Francium
Li Na K Rb Cs Fr
Atomic number 3 11 19 37 55 87
Atomic mass (g mol
–1
) 6.94 22.99 39.10 85.47 132.91 (223)
Electronic [He] 2s
1
[Ne] 3s
1
[Ar] 4s
1
[Kr] 5s
1
[Xe] 6s
1
[Rn] 7s
1
configuration
Ionization 520 496 419 403 376 ~375
enthalpy / kJ mol
–1
Hydration –506 –406 –330 –310 –276 –
enthalpy/kJ mol
–1
Metallic 152 186 227 248 265 –
radius / pm
Ionic radius 76 102 138 152 167 (180)
M
+
 / pm
m.p. / K 454 371 336 312 302 –
b.p / K 1615 1156 1032 961 944 –
Density / g cm
–3
0.53 0.97 0.86 1.53 1.90 –
Standard potentials –3.04 –2.714 –2.925 –2.930 –2.927 –
E
0
/ V for (M
+
 / M)
Occurrence in 18* 2.27** 1.84** 78-12* 2-6* ~ 10
–18 
*
lithosphere
†
This property makes caesium and potassium
useful as electrodes in photoelectric cells.
10.1.6  Chemical Properties
The alkali metals are highly reactive due to
their large size and low ionization enthalpy. The
reactivity of these metals increases down the
group.
(i) Reactivity towards air:  The alkali metals
tarnish in dry air due to the formation of
their oxides which in turn react with
moisture to form hydroxides.  They burn
vigorously in oxygen forming oxides.
Lithium forms monoxide, sodium forms
peroxide, the other metals form
superoxides. The superoxide O
2
–
  ion is
stable only in the presence of large cations
such as K, Rb, Cs.
22
4Li O 2Li O (oxide) +?
222
2Na O Na O (peroxide) +?
22
M O MO (superoxide) +?
(M = K, Rb, Cs)
In all these oxides the oxidation state of the
alkali metal is +1. Lithium shows exceptional
behaviour in reacting directly with nitrogen of
air to form the nitride, Li
3
N as well. Because of
their high reactivity towards air and water,
alkali metals are normally kept in kerosene oil.
Problem 10.1
What is the oxidation state of K in KO
2
?
Solution
The superoxide species is represented as
O
2
–
; since the compound is neutral,
therefore, the oxidation state of potassium
is +1.
*ppm (part per million), ** percentage by weight; † Lithosphere: The Earth’s outer layer: its crust
and part of the upper mantle
Table 10.1  Atomic and Physical Properties of the Alkali Metals
© NCERT
not to be republished
Page 4


291       C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd  Reprint   27.7.6
291 THE s-BLOCK ELEMENTS
The s-block elements of the Periodic Table are those in
which the last electron enters the outermost s-orbital. As
the s-orbital can accommodate only two electrons, two
groups (1 & 2) belong to the s-block of the Periodic Table.
Group 1 of the Periodic Table consists of the elements:
lithium, sodium, potassium, rubidium, caesium and
francium. They are collectively known as the alkali metals.
These are so called because they form hydroxides on
reaction with water which are strongly alkaline in nature.
The elements of Group 2 include beryllium, magnesium,
calcium, strontium, barium and radium. These elements
with the exception of beryllium are commonly known as
the alkaline earth metals. These are so called because their
oxides and hydroxides are alkaline in nature and these
metal oxides are found in the earth’s crust*.
Among the alkali metals sodium and potassium are
abundant and lithium, rubidium and caesium have much
lower abundances (Table 10.1). Francium is highly
radioactive; its longest-lived isotope 
223
Fr has a half-life of
only 21 minutes. Of the alkaline earth metals calcium and
magnesium rank fifth and sixth in abundance respectively
in the earth’s crust. Strontium and barium have much
lower abundances. Beryllium is rare and radium is the
rarest of all comprising only 10
–10
 per cent of igneous
rocks
†
 (Table 10.2, page 299).
The general electronic configuration of s-block elements
is [noble gas]ns
1
 for alkali metals and [noble gas] ns
2
 for
alkaline earth metals.
UNIT 10
After studying this unit, you will be
able to
• • • • • describe the general charact-
eristics of the alkali metals and
their compounds;
• • • • • explain the general characteristics
of the alkaline earth metals and
their compounds;
• • • • • describe the manufacture,
properties and uses of industrially
important sodium and calcium
compounds including Portland
cement;
• • • • • appreciate the biological
significance of sodium,
potassium, magnesium and
calcium.
THE s -BLOCK ELEMENTS
* The thin, rocky outer layer of the Earth is crust.   † A type of rock formed
from magma (molten rock) that has cooled and hardened.
The first element of alkali and alkaline earth metals differs
in many respects from the other members of the group
© NCERT
not to be republished
292       C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd  Reprint   27.7.6
292 CHEMISTRY
Lithium and beryllium, the first elements
of Group 1 and Group 2 respectively exhibit
some properties which are different from those
of the other members of the respective group.
In these anomalous properties they resemble
the second element of the following group.
Thus, lithium shows similarities to magnesium
and beryllium to aluminium in many of their
properties. This type of diagonal similarity is
commonly referred to as diagonal relationship
in the periodic table. The diagonal relationship
is due to the similarity in ionic sizes and /or
charge/radius ratio of the elements.
Monovalent sodium and potassium ions and
divalent magnesium and calcium ions are
found in large proportions in biological fluids.
These ions perform important biological
functions such as maintenance of ion balance
and nerve impulse conduction.
10.1 GROUP 1 ELEMENTS: ALKALI
METALS
The alkali metals show regular trends in their
physical and chemical properties with the
increasing atomic number. The atomic,
physical and chemical properties of alkali
metals are discussed below.
10.1.1 Electronic Configuration
All the alkali metals have one valence electron,
ns
1
 (Table 10.1) outside the noble gas core.
The loosely held s-electron in the outermost
valence shell of these elements makes them the
most electropositive metals. They readily lose
electron to give monovalent M
+
 ions. Hence they
are never found in free state in nature.
increase in atomic number, the atom becomes
larger. The monovalent ions (M
+
) are smaller
than the parent atom. The atomic and ionic
radii of alkali metals increase on moving down
the group i.e., they increase in size while going
from Li to Cs.
10.1.3 Ionization Enthalpy
The ionization enthalpies of the alkali metals
are considerably low and decrease down the
group from Li to Cs. This is because the effect
of increasing size outweighs the increasing
nuclear charge, and the outermost electron is
very well screened from the nuclear charge.
10.1.4 Hydration Enthalpy
The hydration enthalpies of alkali metal ions
decrease with increase in ionic sizes.
Li
+
> Na
+
 > K
+
 > Rb
+
 > Cs
+
Li
+
 has maximum degree of hydration and
for this reason lithium salts are mostly
hydrated, e.g., LiCl· 2H
2
O
10.1.5 Physical Properties
All the alkali metals are silvery white, soft and
light metals. Because of the large size, these
elements have low density which increases
down the group from Li to Cs. However,
potassium is lighter than sodium. The melting
and boiling points of the alkali metals are low
indicating weak metallic bonding due to the
presence of only a single valence electron in
them. The alkali metals and their salts impart
characteristic colour to an oxidizing flame. This
is because the heat from the flame excites the
outermost orbital electron to a higher energy
level. When the excited electron comes back to
the ground state, there is emission of radiation
in the visible region as given below:
Alkali metals can therefore, be detected by
the respective flame tests and can be
determined by flame photometry or atomic
absorption spectroscopy. These elements when
irradiated with light, the light energy absorbed
may be sufficient to make an atom lose electron.
Element Symbol Electronic configuration
Lithium Li 1s
2
2s
1
Sodium Na 1s
2
2s
2
2p
6
3s
1
Potassium K 1s
2
2s
2
2p
6
3s
2
3p
6
4s
1
Rubidium Rb 1s
2
2s
2
2p
6
3s
2
3p
6
3d
10
4s
2
4p
6
5s
1
Caesium Cs 1s
2
2s
2
2p
6
3s
2
3p
6
3d
10
4s
2
4p
6
4d
10
5s
2
5p
6
6s
1 
or [Xe] 6s
1
Francium Fr [Rn]7s
1
10.1.2 Atomic and Ionic Radii
The alkali metal atoms have the largest sizes
in a particular period of the periodic table. With
Metal Li Na K Rb Cs
Colour Crimson Yellow Violet Red Blue
red violet
?/nm 670.8 589.2 766.5 780.0 455.5
© NCERT
not to be republished
293       C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd  Reprint   27.7.6
293 THE s-BLOCK ELEMENTS
Property Lithium Sodium Potassium Rubidium Caesium Francium
Li Na K Rb Cs Fr
Atomic number 3 11 19 37 55 87
Atomic mass (g mol
–1
) 6.94 22.99 39.10 85.47 132.91 (223)
Electronic [He] 2s
1
[Ne] 3s
1
[Ar] 4s
1
[Kr] 5s
1
[Xe] 6s
1
[Rn] 7s
1
configuration
Ionization 520 496 419 403 376 ~375
enthalpy / kJ mol
–1
Hydration –506 –406 –330 –310 –276 –
enthalpy/kJ mol
–1
Metallic 152 186 227 248 265 –
radius / pm
Ionic radius 76 102 138 152 167 (180)
M
+
 / pm
m.p. / K 454 371 336 312 302 –
b.p / K 1615 1156 1032 961 944 –
Density / g cm
–3
0.53 0.97 0.86 1.53 1.90 –
Standard potentials –3.04 –2.714 –2.925 –2.930 –2.927 –
E
0
/ V for (M
+
 / M)
Occurrence in 18* 2.27** 1.84** 78-12* 2-6* ~ 10
–18 
*
lithosphere
†
This property makes caesium and potassium
useful as electrodes in photoelectric cells.
10.1.6  Chemical Properties
The alkali metals are highly reactive due to
their large size and low ionization enthalpy. The
reactivity of these metals increases down the
group.
(i) Reactivity towards air:  The alkali metals
tarnish in dry air due to the formation of
their oxides which in turn react with
moisture to form hydroxides.  They burn
vigorously in oxygen forming oxides.
Lithium forms monoxide, sodium forms
peroxide, the other metals form
superoxides. The superoxide O
2
–
  ion is
stable only in the presence of large cations
such as K, Rb, Cs.
22
4Li O 2Li O (oxide) +?
222
2Na O Na O (peroxide) +?
22
M O MO (superoxide) +?
(M = K, Rb, Cs)
In all these oxides the oxidation state of the
alkali metal is +1. Lithium shows exceptional
behaviour in reacting directly with nitrogen of
air to form the nitride, Li
3
N as well. Because of
their high reactivity towards air and water,
alkali metals are normally kept in kerosene oil.
Problem 10.1
What is the oxidation state of K in KO
2
?
Solution
The superoxide species is represented as
O
2
–
; since the compound is neutral,
therefore, the oxidation state of potassium
is +1.
*ppm (part per million), ** percentage by weight; † Lithosphere: The Earth’s outer layer: its crust
and part of the upper mantle
Table 10.1  Atomic and Physical Properties of the Alkali Metals
© NCERT
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294       C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd  Reprint   27.7.6
294 CHEMISTRY
(ii) Reactivity towards water: The alkali
metals react with water to form hydroxide
and dihydrogen.
22
2M 2H O 2M 2OH H
+-
+? + +
(M = an alkali metal)
It may be noted that although lithium has
most negative E
0
 value (Table 10.1), its
reaction with water is less vigorous than
that of sodium which has the least negative
E
0
 value among the alkali metals. This
behaviour of lithium is attributed to its
small size and very high hydration energy.
Other metals of the group react explosively
with water.
They also react with proton donors such
as alcohol, gaseous ammonia and alkynes.
(iii)Reactivity towards dihydrogen: The
alkali metals react with dihydrogen at
about 673K (lithium at 1073K) to form
hydrides. All the alkali metal hydrides are
ionic solids with high melting points.
2
2M H 2M H
+-
+?
(iv) Reactivity towards halogens : The alkali
metals readily react vigorously with
halogens to form ionic halides, M
+
X
–
.
However, lithium halides are somewhat
covalent. It is because of the high
polarisation capability of lithium ion (The
distortion of electron cloud of the anion by
the cation is called polarisation). The Li
+
 ion
is very small in size and has high tendency
to distort electron cloud around the
negative halide ion. Since anion with large
size can be easily distorted, among halides,
lithium iodide is the most covalent in
nature.
(v) Reducing nature: The alkali metals are
strong reducing agents, lithium being the
most and sodium the least powerful
(Table 10.1). The standard electrode
potential (E
0
) which measures the reducing
power represents the overall change :
2
M(s) M(g) sublimationenthalpy
M(g) M (g) e ionizationenthalpy
M (g) H O M (aq) hydrationenthalpy
+-
++
?
?+
+?
With the small size of its ion, lithium has
the highest hydration enthalpy which
accounts for its high negative E
0
 value and
its high reducing power.
Problem 10.2
The E
0
 for Cl
2
/Cl
– 
is +1.36, for I
2
/I
–
 
 
is
+ 0.53, for Ag
+
 /Ag is  +0.79, Na
+
 /Na is
–2.71 and for Li
+
 /Li is – 3.04. Arrange
the following ionic species in decreasing
order of reducing strength:
I
–
, Ag, Cl
–
, Li, Na
Solution
The order is Li > Na > I
– 
> Ag 
 
> Cl
–
(vi) Solutions in liquid ammonia: The alkali
metals dissolve in liquid ammonia giving
deep blue solutions which are conducting
in nature.
      
33x 3y
M(x y)NH [M(NH)] [e(NH)]
+-
++ ? +
The blue colour of the solution is due to
the ammoniated electron which absorbs
energy in the visible region of light and thus
imparts blue colour to the solution. The
solutions are paramagnetic and on
standing slowly liberate hydrogen resulting
in the formation of amide.
      
(am) 3 2(am) 2
M e NH (1) MNH ½H (g)
+-
++ ? +
(where ‘am’ denotes solution in ammonia.)
In concentrated solution, the blue colour
changes to bronze colour and becomes
diamagnetic.
10.1.7 Uses
Lithium metal is used to make useful alloys,
for example with lead to make ‘white metal’
bearings for motor engines, with aluminium
to make aircraft parts, and with magnesium
to make armour plates. It is used in
thermonuclear reactions. Lithium is also used
to make electrochemical cells. Sodium is used
to make a Na/Pb alloy needed to make PbEt
4
and PbMe
4
. These organolead compounds were
earlier used as anti-knock additives to petrol,
but nowadays vehicles use lead-free petrol.
Liquid sodium metal is used as a coolant in
fast breeder nuclear reactors. Potassium has
© NCERT
not to be republished
Page 5


291       C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd  Reprint   27.7.6
291 THE s-BLOCK ELEMENTS
The s-block elements of the Periodic Table are those in
which the last electron enters the outermost s-orbital. As
the s-orbital can accommodate only two electrons, two
groups (1 & 2) belong to the s-block of the Periodic Table.
Group 1 of the Periodic Table consists of the elements:
lithium, sodium, potassium, rubidium, caesium and
francium. They are collectively known as the alkali metals.
These are so called because they form hydroxides on
reaction with water which are strongly alkaline in nature.
The elements of Group 2 include beryllium, magnesium,
calcium, strontium, barium and radium. These elements
with the exception of beryllium are commonly known as
the alkaline earth metals. These are so called because their
oxides and hydroxides are alkaline in nature and these
metal oxides are found in the earth’s crust*.
Among the alkali metals sodium and potassium are
abundant and lithium, rubidium and caesium have much
lower abundances (Table 10.1). Francium is highly
radioactive; its longest-lived isotope 
223
Fr has a half-life of
only 21 minutes. Of the alkaline earth metals calcium and
magnesium rank fifth and sixth in abundance respectively
in the earth’s crust. Strontium and barium have much
lower abundances. Beryllium is rare and radium is the
rarest of all comprising only 10
–10
 per cent of igneous
rocks
†
 (Table 10.2, page 299).
The general electronic configuration of s-block elements
is [noble gas]ns
1
 for alkali metals and [noble gas] ns
2
 for
alkaline earth metals.
UNIT 10
After studying this unit, you will be
able to
• • • • • describe the general charact-
eristics of the alkali metals and
their compounds;
• • • • • explain the general characteristics
of the alkaline earth metals and
their compounds;
• • • • • describe the manufacture,
properties and uses of industrially
important sodium and calcium
compounds including Portland
cement;
• • • • • appreciate the biological
significance of sodium,
potassium, magnesium and
calcium.
THE s -BLOCK ELEMENTS
* The thin, rocky outer layer of the Earth is crust.   † A type of rock formed
from magma (molten rock) that has cooled and hardened.
The first element of alkali and alkaline earth metals differs
in many respects from the other members of the group
© NCERT
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292       C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd  Reprint   27.7.6
292 CHEMISTRY
Lithium and beryllium, the first elements
of Group 1 and Group 2 respectively exhibit
some properties which are different from those
of the other members of the respective group.
In these anomalous properties they resemble
the second element of the following group.
Thus, lithium shows similarities to magnesium
and beryllium to aluminium in many of their
properties. This type of diagonal similarity is
commonly referred to as diagonal relationship
in the periodic table. The diagonal relationship
is due to the similarity in ionic sizes and /or
charge/radius ratio of the elements.
Monovalent sodium and potassium ions and
divalent magnesium and calcium ions are
found in large proportions in biological fluids.
These ions perform important biological
functions such as maintenance of ion balance
and nerve impulse conduction.
10.1 GROUP 1 ELEMENTS: ALKALI
METALS
The alkali metals show regular trends in their
physical and chemical properties with the
increasing atomic number. The atomic,
physical and chemical properties of alkali
metals are discussed below.
10.1.1 Electronic Configuration
All the alkali metals have one valence electron,
ns
1
 (Table 10.1) outside the noble gas core.
The loosely held s-electron in the outermost
valence shell of these elements makes them the
most electropositive metals. They readily lose
electron to give monovalent M
+
 ions. Hence they
are never found in free state in nature.
increase in atomic number, the atom becomes
larger. The monovalent ions (M
+
) are smaller
than the parent atom. The atomic and ionic
radii of alkali metals increase on moving down
the group i.e., they increase in size while going
from Li to Cs.
10.1.3 Ionization Enthalpy
The ionization enthalpies of the alkali metals
are considerably low and decrease down the
group from Li to Cs. This is because the effect
of increasing size outweighs the increasing
nuclear charge, and the outermost electron is
very well screened from the nuclear charge.
10.1.4 Hydration Enthalpy
The hydration enthalpies of alkali metal ions
decrease with increase in ionic sizes.
Li
+
> Na
+
 > K
+
 > Rb
+
 > Cs
+
Li
+
 has maximum degree of hydration and
for this reason lithium salts are mostly
hydrated, e.g., LiCl· 2H
2
O
10.1.5 Physical Properties
All the alkali metals are silvery white, soft and
light metals. Because of the large size, these
elements have low density which increases
down the group from Li to Cs. However,
potassium is lighter than sodium. The melting
and boiling points of the alkali metals are low
indicating weak metallic bonding due to the
presence of only a single valence electron in
them. The alkali metals and their salts impart
characteristic colour to an oxidizing flame. This
is because the heat from the flame excites the
outermost orbital electron to a higher energy
level. When the excited electron comes back to
the ground state, there is emission of radiation
in the visible region as given below:
Alkali metals can therefore, be detected by
the respective flame tests and can be
determined by flame photometry or atomic
absorption spectroscopy. These elements when
irradiated with light, the light energy absorbed
may be sufficient to make an atom lose electron.
Element Symbol Electronic configuration
Lithium Li 1s
2
2s
1
Sodium Na 1s
2
2s
2
2p
6
3s
1
Potassium K 1s
2
2s
2
2p
6
3s
2
3p
6
4s
1
Rubidium Rb 1s
2
2s
2
2p
6
3s
2
3p
6
3d
10
4s
2
4p
6
5s
1
Caesium Cs 1s
2
2s
2
2p
6
3s
2
3p
6
3d
10
4s
2
4p
6
4d
10
5s
2
5p
6
6s
1 
or [Xe] 6s
1
Francium Fr [Rn]7s
1
10.1.2 Atomic and Ionic Radii
The alkali metal atoms have the largest sizes
in a particular period of the periodic table. With
Metal Li Na K Rb Cs
Colour Crimson Yellow Violet Red Blue
red violet
?/nm 670.8 589.2 766.5 780.0 455.5
© NCERT
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293       C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd  Reprint   27.7.6
293 THE s-BLOCK ELEMENTS
Property Lithium Sodium Potassium Rubidium Caesium Francium
Li Na K Rb Cs Fr
Atomic number 3 11 19 37 55 87
Atomic mass (g mol
–1
) 6.94 22.99 39.10 85.47 132.91 (223)
Electronic [He] 2s
1
[Ne] 3s
1
[Ar] 4s
1
[Kr] 5s
1
[Xe] 6s
1
[Rn] 7s
1
configuration
Ionization 520 496 419 403 376 ~375
enthalpy / kJ mol
–1
Hydration –506 –406 –330 –310 –276 –
enthalpy/kJ mol
–1
Metallic 152 186 227 248 265 –
radius / pm
Ionic radius 76 102 138 152 167 (180)
M
+
 / pm
m.p. / K 454 371 336 312 302 –
b.p / K 1615 1156 1032 961 944 –
Density / g cm
–3
0.53 0.97 0.86 1.53 1.90 –
Standard potentials –3.04 –2.714 –2.925 –2.930 –2.927 –
E
0
/ V for (M
+
 / M)
Occurrence in 18* 2.27** 1.84** 78-12* 2-6* ~ 10
–18 
*
lithosphere
†
This property makes caesium and potassium
useful as electrodes in photoelectric cells.
10.1.6  Chemical Properties
The alkali metals are highly reactive due to
their large size and low ionization enthalpy. The
reactivity of these metals increases down the
group.
(i) Reactivity towards air:  The alkali metals
tarnish in dry air due to the formation of
their oxides which in turn react with
moisture to form hydroxides.  They burn
vigorously in oxygen forming oxides.
Lithium forms monoxide, sodium forms
peroxide, the other metals form
superoxides. The superoxide O
2
–
  ion is
stable only in the presence of large cations
such as K, Rb, Cs.
22
4Li O 2Li O (oxide) +?
222
2Na O Na O (peroxide) +?
22
M O MO (superoxide) +?
(M = K, Rb, Cs)
In all these oxides the oxidation state of the
alkali metal is +1. Lithium shows exceptional
behaviour in reacting directly with nitrogen of
air to form the nitride, Li
3
N as well. Because of
their high reactivity towards air and water,
alkali metals are normally kept in kerosene oil.
Problem 10.1
What is the oxidation state of K in KO
2
?
Solution
The superoxide species is represented as
O
2
–
; since the compound is neutral,
therefore, the oxidation state of potassium
is +1.
*ppm (part per million), ** percentage by weight; † Lithosphere: The Earth’s outer layer: its crust
and part of the upper mantle
Table 10.1  Atomic and Physical Properties of the Alkali Metals
© NCERT
not to be republished
294       C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd  Reprint   27.7.6
294 CHEMISTRY
(ii) Reactivity towards water: The alkali
metals react with water to form hydroxide
and dihydrogen.
22
2M 2H O 2M 2OH H
+-
+? + +
(M = an alkali metal)
It may be noted that although lithium has
most negative E
0
 value (Table 10.1), its
reaction with water is less vigorous than
that of sodium which has the least negative
E
0
 value among the alkali metals. This
behaviour of lithium is attributed to its
small size and very high hydration energy.
Other metals of the group react explosively
with water.
They also react with proton donors such
as alcohol, gaseous ammonia and alkynes.
(iii)Reactivity towards dihydrogen: The
alkali metals react with dihydrogen at
about 673K (lithium at 1073K) to form
hydrides. All the alkali metal hydrides are
ionic solids with high melting points.
2
2M H 2M H
+-
+?
(iv) Reactivity towards halogens : The alkali
metals readily react vigorously with
halogens to form ionic halides, M
+
X
–
.
However, lithium halides are somewhat
covalent. It is because of the high
polarisation capability of lithium ion (The
distortion of electron cloud of the anion by
the cation is called polarisation). The Li
+
 ion
is very small in size and has high tendency
to distort electron cloud around the
negative halide ion. Since anion with large
size can be easily distorted, among halides,
lithium iodide is the most covalent in
nature.
(v) Reducing nature: The alkali metals are
strong reducing agents, lithium being the
most and sodium the least powerful
(Table 10.1). The standard electrode
potential (E
0
) which measures the reducing
power represents the overall change :
2
M(s) M(g) sublimationenthalpy
M(g) M (g) e ionizationenthalpy
M (g) H O M (aq) hydrationenthalpy
+-
++
?
?+
+?
With the small size of its ion, lithium has
the highest hydration enthalpy which
accounts for its high negative E
0
 value and
its high reducing power.
Problem 10.2
The E
0
 for Cl
2
/Cl
– 
is +1.36, for I
2
/I
–
 
 
is
+ 0.53, for Ag
+
 /Ag is  +0.79, Na
+
 /Na is
–2.71 and for Li
+
 /Li is – 3.04. Arrange
the following ionic species in decreasing
order of reducing strength:
I
–
, Ag, Cl
–
, Li, Na
Solution
The order is Li > Na > I
– 
> Ag 
 
> Cl
–
(vi) Solutions in liquid ammonia: The alkali
metals dissolve in liquid ammonia giving
deep blue solutions which are conducting
in nature.
      
33x 3y
M(x y)NH [M(NH)] [e(NH)]
+-
++ ? +
The blue colour of the solution is due to
the ammoniated electron which absorbs
energy in the visible region of light and thus
imparts blue colour to the solution. The
solutions are paramagnetic and on
standing slowly liberate hydrogen resulting
in the formation of amide.
      
(am) 3 2(am) 2
M e NH (1) MNH ½H (g)
+-
++ ? +
(where ‘am’ denotes solution in ammonia.)
In concentrated solution, the blue colour
changes to bronze colour and becomes
diamagnetic.
10.1.7 Uses
Lithium metal is used to make useful alloys,
for example with lead to make ‘white metal’
bearings for motor engines, with aluminium
to make aircraft parts, and with magnesium
to make armour plates. It is used in
thermonuclear reactions. Lithium is also used
to make electrochemical cells. Sodium is used
to make a Na/Pb alloy needed to make PbEt
4
and PbMe
4
. These organolead compounds were
earlier used as anti-knock additives to petrol,
but nowadays vehicles use lead-free petrol.
Liquid sodium metal is used as a coolant in
fast breeder nuclear reactors. Potassium has
© NCERT
not to be republished
295       C:\ChemistryXI\Unit-10\Unit-10-Lay-3(reprint).pmd  Reprint   27.7.6
295 THE s-BLOCK ELEMENTS
a vital role in biological systems. Potassium
chloride is used as a fertilizer. Potassium
hydroxide is used in the manufacture of soft
soap. It is also used as an excellent absorbent
of carbon dioxide. Caesium is used in devising
photoelectric cells.
10.2 GENERAL CHARACTERISTICS OF
THE COMPOUNDS OF THE ALKALI
METALS
All the common compounds of the alkali metals
are generally ionic in nature. General
characteristics of some of their compounds are
discussed here.
10.2.1 Oxides and Hydroxides
On combustion in excess of air, lithium forms
mainly the oxide, Li
2
O (plus some peroxide
Li
2
O
2
), sodium forms the peroxide, Na
2
O
2
 (and
some superoxide NaO
2
) whilst potassium,
rubidium and caesium form the superoxides,
MO
2
. Under appropriate conditions pure
compounds M
2
O, M
2
O
2
 and MO
2
 may be
prepared. The increasing stability of the
peroxide or superoxide, as the size of the metal
ion increases, is due to the stabilisation of large
anions by larger cations through lattice energy
effects. These oxides are easily hydrolysed by
water to form the hydroxides according to the
following reactions :
–
22
MO HO 2M 2OH
+
+? +
–
22 2 22
MO 2HO 2M 2OH HO
+
+? + +
–
22 22 2
2MO 2H O 2M 2OH H O O
+
+? + + +
The oxides and the peroxides are colourless
when pure, but the superoxides are yellow or
orange in colour. The superoxides are also
paramagnetic. Sodium peroxide is widely used
as an oxidising agent in inorganic chemistry.
Problem 10.3
Why is KO
2
 paramagnetic ?
Solution
The superoxide O
2
–
 is paramagnetic
because of  one unpaired electron in p*2p
molecular orbital.
The hydroxides which are obtained by the
reaction of the oxides with water are all white
crystalline solids. The alkali metal hydroxides
are the strongest of all bases and dissolve freely
in water with evolution of much heat on
account of intense hydration.
10.2.2 Halides
The alkali metal halides, MX, (X=F,Cl,Br,I) are
all high melting, colourless crystalline solids.
They can be prepared by the reaction of the
appropriate oxide, hydroxide or carbonate with
aqueous hydrohalic acid (HX). All of these
halides have high negative enthalpies of
formation; the ?
f 
H
0
 values for fluorides
become less negative as we go down the group,
whilst the reverse is true for ?
f 
H
0 
for chlorides,
bromides and iodides. For a given metal
?
f 
H
0 
always becomes less negative from
fluoride to iodide.
The melting and boiling points always
follow the trend: fluoride > chloride > bromide
> iodide. All these halides are soluble in water.
The low solubility of LiF in water is due to its
high lattice enthalpy whereas the low solubility
of CsI is due to smaller hydration enthalpy of
its two ions. Other halides of lithium are soluble
in ethanol, acetone and ethylacetate; LiCl is
soluble in pyridine also.
10.2.3 Salts of Oxo-Acids
Oxo-acids are those in which the acidic proton
is on a hydroxyl group with an oxo group
attached to the same atom e.g., carbonic acid,
H
2
CO
3
 (OC(OH)
2
; sulphuric acid, H
2
SO
4
(O
2
S(OH)
2
). The alkali metals form salts with
all the oxo-acids. They are generally soluble
in water and thermally stable. Their
carbonates (M
2
CO
3
) and in most cases the
hydrogencarbonates (MHCO
3
) also are highly
stable to heat. As the electropositive character
increases down the group, the stability of the
carbonates and hydorgencarbonates increases.
Lithium carbonate is not so stable to heat;
lithium being very small in size polarises a
large CO
3
2–
 ion leading to the formation of more
stable Li
2
O and CO
2
. Its hydrogencarbonate
does not exist as a solid.
© NCERT
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