299THE 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
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300 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 of
the spectrum 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
λ/nm 670.8 589.2 766.5 780.0 455.5
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301THE 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
/ 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.
2 2
4 Li O 2 Li O (oxide)
+→
2 22
2 Na O Na O (peroxide)
+→
2 2
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
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302 CHEMISTRY
(ii) Reactivity towards water: The alkali
metals react with water to form hydroxide
and dihydrogen.
2 2
2M 2H O 2M 2OH H
+
+ →+ +
(M = an alkali metal)
It may be noted that although lithium has
most negative E
value (Table 10.1), its
reaction with water is less vigorous than
that of sodium which has the least negative
E
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
) which measures the reducing
power represents the overall change :
2
M(s) M(g) sublimationenthalpy
M(g) M (g) e ionization enthalpy
M (g) H O M (aq) hydration enthalpy
+−
++
→+
+→
With the small size of its ion, lithium has
the highest hydration enthalpy which
accounts for its high negative E
value and
its high reducing power.
Problem 10.2
The E
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.
3 3x 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
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303THE 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
M O H O 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 π*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
values for fluorides
become less negative as we go down the group,
whilst the reverse is true for
f
H

for chlorides,
bromides and iodides. For a given metal
f
H

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.
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304 CHEMISTRY
10.3 ANOMALOUS PROPERTIES OF
LITHIUM
The anomalous behaviour of lithium is due to
the : (i) exceptionally small size of its atom and
ion, and (ii) high polarising power (i.e., charge/
radius ratio). As a result, there is increased
covalent character of lithium compounds which
is responsible for their solubility in organic
solvents. Further, lithium shows diagonal
relationship to magnesium which has been
discussed subsequently.
10.3.1 Points of Difference between
Lithium and other Alkali Metals
(i) Lithium is much harder. Its m.p. and b.p.
are higher than the other alkali metals.
(ii) Lithium is least reactive but the strongest
reducing agent among all the alkali metals.
On combustion in air it forms mainly
monoxide, Li
2
O and the nitride, Li
3
N unlike
other alkali metals.
(iii) LiCl is deliquescent and crystallises as a
hydrate, LiCl.2H
2
O whereas other alkali
metal chlorides do not form hydrates.
(iv) Lithium hydrogencarbonate is not
obtained in the solid form while all other
elements form solid hydrogencarbonates.
(v) Lithium unlike other alkali metals forms
no ethynide on reaction with ethyne.
(vi) Lithium nitrate when heated gives lithium
oxide, Li
2
O, whereas other alkali metal
nitrates decompose to give the
corresponding nitrite.
3 2 22
4LiNO 2 Li O 4 NO O
++
2
3 22
NaNO 2 NaNO O
→+
(vii) LiF and Li
2
O are comparatively much less
soluble in water than the corresponding
compounds of other alkali metals.
10.3.2 Points of Similarities between
Lithium and Magnesium
The similarity between lithium and magnesium
is particularly striking and arises because of
their similar sizes : atomic radii, Li = 152 pm,
Mg = 160 pm; ionic radii : Li
+
= 76 pm,
Mg
2+
= 72 pm. The main points of similarity are:
(i) Both lithium and magnesium are harder
and lighter than other elements in the
respective groups.
(ii) Lithium and magnesium react slowly with
water. Their oxides and hydroxides are
much less soluble and their hydroxides
decompose on heating. Both form a nitride,
Li
3
N and Mg
3
N
2
, by direct combination
with nitrogen.
(iii) The oxides, Li
2
O and MgO do not combine
with excess oxygen to give any superoxide.
(iv) The carbonates of lithium and magnesium
decompose easily on heating to
form the oxides and CO
2
. Solid
hydrogencarbonates are not formed by
lithium and magnesium.
(v) Both LiCl and MgCl
2
are soluble in ethanol.
(vi) Both LiCl and MgCl
2
are deliquescent and
crystallise from aqueous solution as
hydrates, LiCl·2H
2
O and MgCl
2
·8H
2
O.
10.4 SOME IMPORTANT COMPOUNDS OF
SODIUM
Industrially important compounds of sodium
include sodium carbonate, sodium hydroxide,
sodium chloride and sodium bicarbonate. The
large scale production of these compounds
and their uses are described below:
Sodium Carbonate (Washing Soda),
Na
2
CO
3
·10H
2
O
Sodium carbonate is generally prepared by
Solvay Process. In this process, advantage is
taken of the low solubility of sodium
hydrogencarbonate whereby it gets
precipitated in the reaction of sodium chloride
with ammonium hydrogencarbonate. The
latter is prepared by passing CO
2
to a
concentrated solution of sodium chloride
saturated with ammonia, where ammonium
carbonate followed by ammonium
hydrogencarbonate are formed. The equations
for the complete process may be written as :
(
)
32 2 43
2
2 NH H O CO NH CO
++
(
)
4 32 2 43
2
NH CO H O CO 2 NH HCO
++
43
4 3
NH HCO NaCl NH Cl NaHCO
+→ +
Sodium hydrogencarbonate crystal
separates. These are heated to give sodium
carbonate.
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305THE s-BLOCK ELEMENTS
3 23 2 2
2 NaHCO Na CO CO H O
++
In this process NH
3
is recovered when the
solution containing NH
4
Cl is treated with
Ca(OH)
2
. Calcium chloride is obtained as a
by-product.
(
)
4 3 22
2
2 NH Cl Ca OH 2 NH CaCl H O
+ →++
It may be mentioned here that Solvay
process cannot be extended to the
manufacture of potassium carbonate because
potassium hydrogencarbonate is too soluble
to be precipitated by the addition of
ammonium hydrogencarbonate to a saturated
solution of potassium chloride.
Properties : Sodium carbonate is a white
crystalline solid which exists as a decahydrate,
Na
2
CO
3
·10H
2
O. This is also called washing
soda. It is readily soluble in water. On heating,
the decahydrate loses its water of crystallisation
to form monohydrate. Above 373K, the
monohydrate becomes completely anhydrous
and changes to a white powder called soda ash.
375 K
23 2 2 32 2
Na CO 10H O Na CO H O 9H O
→ +
373K
232 23 2
Na CO H O Na CO H O
>
 +
Carbonate part of sodium carbonate gets
hydrolysed by water to form an alkaline
solution.
2– ––
32 3
CO H O HCO OH
+→ +
Uses:
(i) It is used in water softening, laundering
and cleaning.
(ii) It is used in the manufacture of glass,
soap, borax and caustic soda.
(iii) It is used in paper, paints and textile
industries.
(iv) It is an important laboratory reagent both
in qualitative and quantitative analysis.
Sodium Chloride, NaCl
The most abundant source of sodium chloride
is sea water which contains 2.7 to 2.9% by
mass of the salt. In tropical countries like India,
common salt is generally obtained by
evaporation of sea water. Approximately 50
lakh tons of salt are produced annually in
India by solar evaporation. Crude sodium
chloride, generally obtained by crystallisation
of brine solution, contains sodium sulphate,
calcium sulphate, calcium chloride and
magnesium chloride as impurities. Calcium
chloride, CaCl
2
, and magnesium chloride,
MgCl
2
are impurities because they are
deliquescent (absorb moisture easily from the
atmosphere). To obtain pure sodium chloride,
the crude salt is dissolved in minimum amount
of water and filtered to remove insoluble
impurities. The solution is then saturated with
hydrogen chloride gas. Crystals of pure
sodium chloride separate out. Calcium and
magnesium chloride, being more soluble than
sodium chloride, remain in solution.
Sodium chloride melts at 1081K. It has a
solubility of 36.0 g in 100 g of water at 273 K.
The solubility does not increase appreciably
with increase in temperature.
Uses :
(i) It is used as a common salt or table salt for
domestic purpose.
(ii) It is used for the preparation of Na
2
O
2
,
NaOH and Na
2
CO
3
.
Sodium Hydroxide (Caustic Soda), NaOH
Sodium hydroxide is generally prepared
commercially by the electrolysis of sodium
chloride in Castner-Kellner cell. A brine
solution is electrolysed using a mercury
cathode and a carbon anode. Sodium metal
discharged at the cathode combines with
mercury to form sodium amalgam. Chlorine
gas is evolved at the anode.
Hg
Cathode : Na e Na amalgam
+−
+ →
2
1
Anode : Cl Cl e
2
→+
The amalgam is treated with water to give
sodium hydroxide and hydrogen gas.
2Na-amalgam + 2H
2
O2NaOH+ 2Hg +H
2
Sodium hydroxide is a white, translucent
solid. It melts at 591 K. It is readily soluble in
water to give a strong alkaline solution.
Crystals of sodium hydroxide are deliquescent.
The sodium hydroxide solution at the surface
reacts with the CO
2
in the atmosphere to form
Na
2
CO
3
.
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306 CHEMISTRY
Uses: It is used in (i) the manufacture of soap,
paper, artificial silk and a number of chemicals,
(ii) in petroleum refining, (iii) in the purification
of bauxite, (iv) in the textile industries for
mercerising cotton fabrics, (v) for the
preparation of pure fats and oils, and (vi) as a
laboratory reagent.
Sodium Hydrogencarbonate (Baking
Soda), NaHCO
3
Sodium hydrogencarbonate is known as
baking soda because it decomposes on heating
to generate bubbles of carbon dioxide (leaving
holes in cakes or pastries and making them
light and fluffy).
Sodium hydrogencarbonate is made by
saturating a solution of sodium carbonate with
carbon dioxide. The white crystalline powder
of sodium hydrogencarbonate, being less
soluble, gets separated out.
23 2 2 3
Na CO H O CO 2 NaHCO
++
Sodium hydrogencarbonate is a mild
antiseptic for skin infections. It is used in fire
extinguishers.
10.5 BIOLOGICAL IMPORTANCE OF
SODIUM AND POTASSIUM
A typical 70 kg man contains about 90 g of Na
and 170 g of K compared with only 5 g of iron
and 0.06 g of copper.
Sodium ions are found primarily on the
outside of cells, being located in blood plasma
and in the interstitial fluid which surrounds
the cells. These ions participate in the
transmission of nerve signals, in regulating the
flow of water across cell membranes and in the
transport of sugars and amino acids into cells.
Sodium and potassium, although so similar
chemically, differ quantitatively in their ability
to penetrate cell membranes, in their transport
mechanisms and in their efficiency to activate
enzymes. Thus, potassium ions are the most
abundant cations within cell fluids, where they
activate many enzymes, participate in the
oxidation of glucose to produce ATP and, with
sodium, are responsible for the transmission
of nerve signals.
There is a very considerable variation in the
concentration of sodium and potassium ions
found on the opposite sides of cell membranes.
As a typical example, in blood plasma, sodium
is present to the extent of 143 mmolL
–1
,
whereas the potassium level is only
5 mmolL
–1
within the red blood cells. These
concentrations change to 10 mmolL
–1
(Na
+
) and
105 mmolL
–1
(K
+
). These ionic gradients
demonstrate that a discriminatory mechanism,
called the sodium-potassium pump, operates
across the cell membranes which consumes
more than one-third of the ATP used by a
resting animal and about 15 kg per 24 h in a
resting human.
10.6 GROUP 2 ELEMENTS : ALKALINE
EARTH METALS
The group 2 elements comprise beryllium,
magnesium, calcium, strontium, barium and
radium. They follow alkali metals in the
periodic table. These (except beryllium) are
known as alkaline earth metals. The first
element beryllium differs from the rest of the
members and shows diagonal relationship to
aluminium. The atomic and physical
properties of the alkaline earth metals are
shown in Table 10.2.
10.6.1 Electronic Configuration
These elements have two electrons in the
s -orbital of the valence shell (Table 10.2). Their
general electronic configuration may be
represented as [noble gas] ns
2
. Like alkali
metals, the compounds of these elements are
also predominantly ionic.
Element Symbol Electronic
configuration
Beryllium Be 1s
2
2s
2
Magnesium Mg 1s
2
2s
2
2p
6
3s
2
Calcium Ca 1s
2
2s
2
2p
6
3s
2
3p
6
4s
2
Strontium Sr 1s
2
2s
2
2p
6
3s
2
3p
6
3d
10
4s
2
4p
6
5s
2
Barium Ba 1s
2
2s
2
2p
6
3s
2
3p
6
3d
10
4s
2
4p
6
4d
10
5s
2
5p
6
6s
2
or
[Xe]6s
2
Radium Ra [Rn]7s
2
10.6.2 Atomic and Ionic Radii
The atomic and ionic radii of the alkaline earth
metals are smaller than those of the
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307THE s-BLOCK ELEMENTS
corresponding alkali metals in the same
periods. This is due to the increased nuclear
charge in these elements. Within the group, the
atomic and ionic radii increase with increase
in atomic number.
10.6.3 Ionization Enthalpies
The alkaline earth metals have low ionization
enthalpies due to fairly large size of the atoms.
Since the atomic size increases down the
group, their ionization enthalpy decreases
(Table 10.2). The first ionisation enthalpies of
the alkaline earth metals are higher than those
of the corresponding Group 1 metals. This is
due to their small size as compared to the
corresponding alkali metals. It is interesting
to note that the second ionisation enthalpies
of the alkaline earth metals are smaller than
those of the corresponding alkali metals.
10.6.4 Hydration Enthalpies
Like alkali metal ions, the hydration enthalpies
of alkaline earth metal ions decrease with
increase in ionic size down the group.
Be
2+
> Mg
2+
> Ca
2+
> Sr
2+
> Ba
2+
The hydration enthalpies of alkaline earth
metal ions are larger than those of alkali metal
ions. Thus, compounds of alkaline earth metals
are more extensively hydrated than those of
alkali metals, e.g., MgCl
2
and CaCl
2
exist as
MgCl
2
.6H
2
O and CaCl
2
· 6H
2
O while NaCl and
KCl do not form such hydrates.
10.6.5 Physical Properties
The alkaline earth metals, in general, are silvery
white, lustrous and relatively soft but harder
than the alkali metals. Beryllium and
magnesium appear to be somewhat greyish.
The melting and boiling points of these metals
are higher than the corresponding alkali metals
due to smaller sizes. The trend is, however, not
systematic. Because of the low ionisation
enthalpies, they are strongly electropositive in
nature. The electropositive character increases
down the group from Be to Ba. Calcium,
Property Beryllium Magnesium Calcium Strontium Barium Radium
Be Mg Ca Sr Ba Ra
Atomic number 4 12 20 38 56 88
Atomic mass (g mol
–1
) 9.01 24.31 40.08 87.62 137.33 226.03
Electronic [He] 2s
2
[Ne] 3s
2
[Ar] 4s
2
[Kr] 5s
2
[Xe] 6s
2
[Rn] 7s
2
configuration
Ionization 899 737 590 549 503 509
enthalpy (I) / kJ mol
–1
Ionization 1757 1450 1145 1064 965 979
enthalpy (II) /kJ mol
–1
Hydration enthalpy 2494 1921 –1577 1443 1305
(kJ/mol)
Metallic 111 160 197 215 222
radius / pm
Ionic radius 31 72 100 118 135 148
M
2+
/ pm
m.p. / K 1560 924 1124 1062 1002 973
b.p / K 2745 1363 1767 1655 2078 (1973)
Density / g cm
–3
1.84 1.74 1.55 2.63 3.59 (5.5)
Standard potential –1.97 –2.36 –2.84 –2.89 2.92 –2.92
E
/ V for (M
2+
/ M)
Occurrence in 2* 2.76** 4.6** 384* 390 * 10
–6
*
lithosphere
Table 10.2 Atomic and Physical Properties of the Alkaline Earth Metals
*ppm (part per million); ** percentage by weight
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