THE s-BLOCK ELEMENTS

               1. Introduction

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 . Francium is highly radioactive; its longest-lived isotope 223Fr 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. The general electronic configuration of s-block elements is [noble gas] n s¹for alkali metals and [noble gas] n s¹ 2 for alkaline earth metals. 

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.

 2. 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.

1. Electronic Configuration : 

All the alkali metals have one valence electron,

n s¹ 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.

2.  Atomic and Ionic Radii

The alkali metal atoms have the largest sizes

in a particular period of the periodic table. With

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 L i to C s.

3. I o n i z a t i o n Enthalpy

The i o n i z a t i o n enthalpy of the alkali metals

are considerably low and decrease down the

group from L i to C s. 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.

4. Hydration Enthalpy

The hydration enthalpy of alkali metal ions

decrease with increase in ionic sizes.

L i+ > Na+ > K+  > R b+  > C s+ . 

L i+  has maximum degree of hydration and 

for this reason lithium salts are mostly

hydrated, e.g., L i Cl· 2H2O

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 L i to C s. 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. 

6. Chemical Properties

The alkali metals are highly reactive due to

their large size and low i o n i z a t i o n 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 

super oxides. The superoxide O2 –  ion is 

stable only in the presence of large cation

such as K, R b, C s . 4Li O 2Li O (oxide)

2Na O Na O (peroxide) + →2 22

M O MO (superoxide) + →2 2

(M = K, R b, C s)

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, L i 3N as well. Because of

their high reactivity towards air and water,

alkali metals are normally kept in kerosene oil.

(i i) Reactivity towards water: The alkali metals react with water to form hydroxide and dihydrogen. 

2M+2H2O→2M+2OH + H2

(M = an alkali metal)

It may be noted that although lithium has most negative E value, 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 al k y n es.

 

(i i i) 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.

2M+2H2O − + →2M+ H2

(i v) 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 L i+  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 . The standard electrode  potential (E) which measures the reducing power represents the overall change : 2 

M(s) M(g) sublimation enthalpy M(g) M (g) e i o n i z a t i o n enthalpy M (g) H O M (a q) 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.

(vi) Solutions in liquid ammonia: The alkali

metals dissolve in liquid ammonia giving

deep blue solutions which are conducting

in nature. 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. In concentrated solution, the blue colour changes to bronze colour and becomes diamagnetic.

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/P b alloy needed to make P b E t4 and P b M e4. These organ o lead 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

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.

3. GENERAL CHARACTERISTIC 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.

1. Oxides and Hydroxides

On combustion in excess of air, lithium forms

mainly the oxide, L i 2 O (plus some peroxide

L i 2O2), sodium forms the peroxide, N a 2O2 (and some superoxide Na O2) whilst potassium, 

rubidium and caesium form the super oxides,

M O2. Under appropriate conditions pure

compounds M 2O, M2 O2 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 cation through lattice energy

effects. These oxides are easily hydrolysed by

water to form the hydroxides. 

The oxides and the peroxides are colourless

when pure, but the super oxides are yellow or

orange in colour. The super oxides are also

paramagnetic. Sodium peroxide is widely used

as an oxidising agent in inorganic chemistry

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.

2. Halides

The alkali metal halides, MX, (X=F,Cl,B r,I) are

all high melting, colourless crystalline solids.

They can be prepared by the reaction of the

appropriate oxide, hydroxide or carbonate with

aqueous hydro ha l i c acid (HX). All of these

halides have high negative enthalpy 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 s o l u b i l i t y of L iF in water is due to its high lattice enthalpy whereas the low s o lu b i l i ty of  C Si is due to smaller hydration enthalpy of its two ions. Other halides of lithium are soluble in ethanol, acetone and ethyl acetate; L i Cl is soluble in pyridine also. 

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,

H2 CO 3 (O C(OH)2; sulphuric acid, H2 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 (M2 CO 3) and in most cases the

hydrogen carbonates (M H CO 3) also are highly

stable to heat. As the electropositive character

increases down the group, the stability of the

carbonates and hydrogen carbonates 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 L i 2O and CO 2. Its hydrogen carbonate does not exist as a solid.

4. ANOMALOUS PROPERTIES OF LITHIUM

The anomalous behaviour of lithium is due to

the :

 (i) exceptionally small size of its atom and

ion, and

 (i i) high polarising power (i.e., charge/

radius ratio). As a result, there is increased

covalent character of lithium compounds which

is responsible for their s o l u b i l i t y in organic

solvents. Further, lithium shows diagonal

relationship to magnesium which has been

discussed subsequently.

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.

(i i) Lithium is least reactive but the strongest

reducing agent among all the alkali metals.

On combustion in air it forms mainly monoxide, L i 2O and the nitride, L i 3 N unlike other alkali metals.

(i i i) L i Cl is deliquescent and crystallises as a

hydrate, L i Cl.2H2O whereas other alkali

metal chlorides do not form hydrates.

(i v) Lithium hydrogen carbonate is not

obtained in the solid form while all other

elements form solid hydrogen carbonates.

(v) Lithium unlike other alkali metals forms

no e t hy nide on reaction with eth y n e.

(vi) Lithium nitrate when heated gives lithium

oxide, L i 2 O, whereas other alkali metal

nitrates decompose to give the

corresponding nitrite.

(v i i) L iF and L i 2O are comparatively much less

soluble in water than the corresponding compounds of other alkali metals.

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, L i = 152 pm,

Mg = 160 pm; ionic radii : L i+ = 76 pm, M g2+= 72 pm. The main points of similarity are:

(i) Both lithium and magnesium are harder

and lighter than other elements in the

respective groups.

(i i) 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,

L i 3N and Mg 3N2, by direct combination

with nitrogen.

(i i i) The oxide, L i 2 O and Mg O do not combine

with excess oxygen to give any superoxide.

(i v) The carbonates of lithium and magnesium

decompose easily on heating to form the oxides and CO 2. Solid hydrogen carbonates are not formed by lithium and magnesium.

(v) Both L i Cl and Mg Cl 2 are soluble in ethanol.

(vi) Both L i Cl and Mg Cl 2 are deliquescent and

crystallise from aqueous solution as hydrates, L i Cl·2H2O and Mg Cl 2·8H2O.