S-Block Elements (Part -1) Chemistry Notes | EduRev

Inorganic Chemistry

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Chemistry : S-Block Elements (Part -1) Chemistry Notes | EduRev

The document S-Block Elements (Part -1) Chemistry Notes | EduRev is a part of the Chemistry Course Inorganic Chemistry.
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1. Electronic Configuration

S-Block Elements (Part -1) Chemistry Notes | EduRev

The silvery lustre of alkali metals is due to the presence of highly mobile electrons of the metallic lattice. There being only a single electron per atom, the metallic bonding is not so strong. As a result of this, these metals are soft in nature. However, the softness increases with increases of atomic number because there is continuous decrease of metallic bond strength on account of increase in atomic size.

  • Bigger is the size of metal kernel weaker is the metallic bo nding. 

2. Atomic and ionic radii

At radii (Å)

3. Density
All are light metals. The sensities of metals Li, Na and K Are lighter than water. Density gradually increases in moving down from Li to Cs. Potassium is however, lighter than sodium. Sequence of densities: Li < Na < K < Rb < Cs

4. Structures of the metals, hardness and cohesive energy 

  • At normal temperatures all the Group 1 metals adopt a body -centred cubic type of lattice with a coordination number of 8. However, at very low temperatures lithium forms a hexagonal closepacked structure with a coordination number of 12.
  •  The cohesive energy is the force holding the atoms or ions together in the solid.
  •  The atoms become larger on descending the group from lithium to caesium, so the bonds are weaker, the cohesive energy decreases and the softness of the metals increase.
  •  The crystal structures of Li2O, Na2O, K2O and Rb2O are anti fluorite structures, Cs2O has an anti CdCl2 layer structure.

5. Melting and boiling points

  •  The cohesive energy decreases down the group, and the melting points decrease corresponding.

The energy binding the atoms in the crystal lattices of these metals in relatively low on account of a single electron in the valency shell. Consequently, the metals have low melting and boiling points. These decrease in moving down from Li to Cs as the metallic bond strength decreases or cohesive force decrease.

6. Ionisation energies and electropositive character 

Due to their large size, the outermost electron is far from the nucleus and can easily be removed. Their ionization potentials are relatively low. Thus, the metals have a great tendency to lose the ns1 electron to change into M + ions. These metals are highly electropositive in nature. As the ionization potential decreases from Li to Cs, the electropositive character increases, i.e. metallic character increases. The reactivity of these metals increases from Li to Cs.

Ionization Potential (eV)

S-Block Elements (Part -1) Chemistry Notes | EduRev

  • The ns1 electron is so loosely held that even the lo w energy photons (light) can eject this electron from the metal surface. This property is termed as photo electronic effect. K and Cs are used in photoelectric cells which are sensitive to blue light. 

7. Oxidation States 

The alkali metals can lose their ns1 electron quite easily to form univalent positive ion, M+. The ion has a stable configuration of an inert gas. These metals are univalent in nature and show electro valency, i.e., form electro valent compounds.

  •  Since the electron configuration of M + ions are similar to those of inert gases, these ions have no unpaired electrons and consequently are colourless and diamagnetic in nature. 
  •  Hydration of Ions, hydrated radii and hydration energy
    The salts of alkali metals are ionic and soluble in water. The solubility is due to the test that cations get hydrated by water molecules M+ + aq   —→  [M(aq)]+  Hydrated Cation
  • The smaller the cation, the greater is the degree of its hydration of M + ions & decrease from Li to Cs+
    • Consequently the radii of the hydrated ion decreases from Li+ to Cs
    •  The ionic conductance of these hydrated ions increases from [Li(aq)] + to [Cs(aq)]+.
    • Hydration of ions is an exothermic process.
    •  The energy released when one gram mole of an ion is dissolved in water to get it hydrated is called hydration energy. Since the degree of hydration decreases from Li + to Cs+, the hydration energy of alkali metal ion also decreases from Li+ to Cs+.
  • Some water molecules touch the metal ions and bond to it, forming a complex. These water molecules consititute the primary shell of water. Thus Liis tetrahedrally surrounded by four water molecules sphybridization. 
  •  With the heavier ions, particularly Rb+ and Cs+, the number of water molecules increases to six. VSEPR theory predicts an octahedral structure (d2sp3 hybridization). 
  •  A secondary layer of water molecules further hydrates the ions, tho ugh these are o nly held by weak ion-dipole attractive forces. 

9. Conductivity

The alkali metals are good conductors of heat and electricity. This is due to the presence of loosely held valency electrons which are free to move throughout the metal structure.

10. Heat of atomization 

  • Heat of atomization decreases from Li to Cs.
  •  This is due to the decrease in the metallic bond strength from Li to Cs

11. Flame colours & spectra

The alkali metals and their salts impart a characteristic colour to flame.

The colour actually arises from electronic transitions in short-lived species which are formed
                                                Li                     Na                  K                  Rb                   Cs
Momentarily in the flame        Crimson red     Golden red     Pale violet   Violet               Violet

  • The rea son for flame colouration is that the energy of the flame cau ses an excitation of the outermost electrons which on return to their original position give out the energy so absorbed in the visible region. The energy released is minimu m ion the case of Li + and increases from Li to Cs+. Thu s the frequency of the light emitted increases in accordance with the formula E = hν. The frequency of light in lithium is minimum which corresponds to red region of the spectra.
  • Compounds of Group1 metals are typically white, except those where the anion is coloured. For example sodium chromate Na2(CrO4) (yellow), potassium dichromate K2[Cr2O7] (orange), and potassium permanganate K[MnO4] (deep purple).

12. Reducing nature 

An element, which acts as a reducing agent, must have low ionization energy. Alkali metals act as strong reducing agents as their ionization energy values are low. Since ionization energy decreases on moving down from Li to Cs, the reducing property should increases in the same order but from E°ox values it is observed that Li is the strongest reducing agent a mongst alkali metals in solution a s E°ox value of Li is maximum.
It is due to high hydration energy of Li + in a queous solutions.

Oxidation Potential (V)

Chemical Properties

 1. Action with Air

On exposure to moist air, all alkali metals except lithium tarnish quickly. The effect of atmosphere increases from Li to Cs. These are, therefore always kept under kerosene oil to protect them from air. 

S-Block Elements (Part -1) Chemistry Notes | EduRev

  • Lithium when heated in air combines with nitrogen to form nitride, it is due to diagonal relationship with magnesium.

2. Action with Water

Alkali metals decompose water with the evolution of hydrogen gas.

S-Block Elements (Part -1) Chemistry Notes | EduRev

3. Action with hydrogen The alkali metals combine directly with hydrogen to form crystalline hydrides of the formula M. H. These hydrides are ionic and contain the hydrides ion. H
2M +H2 ——→ 2MH
The ionic character of the bonds in these hydride increases from LiH to CsH and the stability decreases in the same order. They are powerful reducing agents especially at high temperatures.
SiCl4 + 4NaH ——→ SiH4 + 4NaCl 

4. Action with Oxygen (Oxides and Hydroxides)

  • Down the group affinity towards oxygen increases.
  • The metals all burn in air to from oxides, thought the produ ct varies depending on the metal. lithium forms the monoxide Li2O (and some peroxide Li2O2), sodiu m forms the peroxide Na2O2 and some monoxide Na2O), and the others form superoxide’s of the type MO2

S-Block Elements (Part -1) Chemistry Notes | EduRev

The normal oxides ‘M2O’ react with water to form hydroxides. 

M2O + H2O ——→ 2MOH
The basic nature of the oxides (M2O) increases gradually on moving down in the group. The hydroxides (MOH) are colourless, strong alkaline and corrosive compounds. These are soluble in water and dissolve with evolution of heat. The hydroxides are thermally stable except LiOH. The relative strength of the hydroxides increases from LiOH to CsOH.
CsOH > RbOH > KOH > NaOH > LiOH

5. Action with halogens The alkali metals directly react with halogens forming the halides of the type MX. 2M + X2 → 2MX
With the exception of certain lithium haloids, the metal halides are ionic compounds (M+X–). The halides are crystalline and have high melting and boiling points. The fused halides are good conductors of electricity and are u sed for the preparation of alkali metals. All halides except LiF dissolve in water.

  •  The alkali metal halides react with the halogens and interhalogen compounds forming ionic polyhalide compounds.

KI + I2 → K[I3],                 KBr + ICI → K[BrICI],                 KF + BrF3 → K[BrF4]

6. Natural of Oxysalts
Alkali metals readily react with oxyacids forming corresponding salts with evolution of hydrogen. Lithium salt behave abnormally due to polarizing power of Li+ ion (small size) and lattice energy effects.

7. Natural of carbonates and bicarbonates
As the electropositive character increases from Li to Cs, the solubility of the carbonates increases in the same order
Cs2CO3 > Rb2CO3 > K2CO3 > Na2CO3 > Li2CO3

Li2CO3 decomposes on heating and is insoluble in water. The aqueous solution of carbonates are alkaline. This is due to hydrolysis as carbonates are salts of strong bases and weak acid (H2CO3) carbonic acid).
S-Block Elements (Part -1) Chemistry Notes | EduRev⇆  S-Block Elements (Part -1) Chemistry Notes | EduRev

The bicarbonates, MHCO3, of the alkali metals, with the exception of lithium, are known in solid state. The bicarbonates are soluble in water. On heating, bicarbonates decompose into carbonates with evolution of CO2.
2MHCO3 ——→ +M2CO3 + H2O + CO2
The abnormal behaviour of Li2CO3 towards heat can be explained in the following manner. The Li ion exerts a strong polarizing action and distorts the electron cloud of the nearby atom the large CO32- ion. This results in the weak ening of the C—O bond and strengthening of the Li—O bond. This ultimately facilitates the decomposition of Li2CO3 into Li2O and CO2. The lattice energy Li2O is higher than the lattice energy of carbonate. This also favours the decomposition of Li2CO3.

  • Lithium due to its less electropositive nature, does not form solid bicarbonate and LiHCO 3 exists in solutions only.

8. Natural of nitrates 

Nitrates of the type, MNO3, are known. These are colourless, soluble in water and electrovalent in nature. The nitrates do not undergo hydrolysis. With the exception of LiNO3, the other nitrates decompose to nitrites and oxygen.
2MNO3 ——→ 2MNO+ O2
Lithium nitrate decomposes to oxide on heating, it is due to diagonal relationship with magnesium.

S-Block Elements (Part -1) Chemistry Notes | EduRev

9. Nature of Sulphates 

Sulphates of the type M2SOare known. With the exception of Li2SO4, other sulphates are solu ble in water. The sulphates when fused with carbon form sulphides.
M2SO+ 4C ——→ M2S + 4CO
The sulphates of alkali metals form double salts with the sulphates of the trivalent metals like Fe, Al, Cr, etc. The double sulphates crystallize with large number of water molecules are potash alum K2SO4, Al2(SO4)3.24H2O consists of 24 water molecules. Sulphate of lithium are not known to form alum.

10. Action of liquid ammonia 

The alkali metals dissolve in liquid ammonia without the evolution of hydrogen. The colour of the dilute solutions is blue. The metal atom loses electron and it combines with ammonia molecule.
M → M+ (in liquid ammonia) + e (ammoniated)
M+(x + y) NH3 → [M(NH3)x]+ + e(NH3)y solvated electron

On heating its blue colour changes to bronze. It is ammoniated electron which is responsible for colour.
The solutions are good conductors of electricity and have strong reducing properties. The solutions are paramagnetic in nature.
When dry ammonia is passed over hot metal, amides are formed.

S-Block Elements (Part -1) Chemistry Notes | EduRev

The amides are decomposed by cold water with evolution of NH3 MNH2 + 2NH3 ——→ MOH + NH3
Recent studies proved the existence of Li(NH3)4, a golden yellow solid.

11. Formation of alloys 

The alkali metals form alloys amongst themselves and with other metals. These combine with mercury readily forming amalgams.

12. Complex Formation

Alkali metals have a very little tendency to form complexes, Lithium being small in size forms certain complexes but this tendency decreases as the size increases.


Lithium shows resemblance with magnesium, an element of group IIA. This resemblance is termied as diagonal relationship.

S-Block Elements (Part -1) Chemistry Notes | EduRev

Reason for the diagonal relationship are the following:

(i) Electronegativities of Li and Mg are quite comparable (Li = 1.00 Mg = 1.20)

(ii) Atomic ra dii and ionic ra dii of Li and Mg are not very mu ch different.
Atomic radii (Å)               Li-1.23,              Mg-1.36
Ionic radii (Å)                  Li+-00.60,          Mg2+-0.65

(iii) Atomic volumes of Li and Mg are go to similar              Li-12.97 mL,              Mg-13.97 mL

(iv) Both have high polarizing power (ionic potential) Polarizing powerS-Block Elements (Part -1) Chemistry Notes | EduRev

Cations with large ionic potentials have a tendency to polarize the anions and to give partial covalent character to compounds.
Lithium resembles magnesium in the following respects.

(a) Both Li and Mg are harder and have higher melting points than the other metals in their respective groups.

(b) Li like Mg decomposes water slowly to liberate hydrogen.

2Li + 2H2O ——→ 2LiOH + H2; Mg + 2H2O ——→ Mg(OH)2 + H

(c) Both elements combine with nitrogen on heating.

6Li + N2 ——→ 2Li3N; 3Mg + N2 ——→ Mg+ N

Both the nitrides are decomposed by water with evolution of ammonia (NH3)

Li3N + 3H2O ——→ 3LiOH + NH3; Mg3N+ 6H2O ——→ 3Mg(OH)2 + 2NH3

(d) Both Li and Mg combine with carbon on heating

2Li + 3H2O ——→ 3Li2C2; Mg + 2C ——→ MgC

Both the carbides yield C2H2 with water.

(e) Lithium forms only monoxide when heated in oxygen.
Mg also forms the monoxide

4Li + O2 ——→ 2Li2O; 2Mg + O ——→ 2MgO

Both the are less soluble in water.

(f) Hydroxides of Li and Mg are weak bases and are slightly soluble in water. Both decomposes on heating

2LiOH ——→ Li2O + H2O; Mg(OH)2 ——→ MgO + H2O

(g) Lithium fluoride, phosphate, oxalate and carbonate lik e the corresponding salts of Mg, are sparingly soluble in water.

(h) Carbonates of Li and Mg decomposes on heating.

Li2CO3 ——→ Li2O + CO2; MgCO ——→ MgO + CO2

(i) Nitrates of Li and Mg decompose on heating giving mixture of nitrogen dioxide and oxygen.

4Li2 NO3 ——→ 2Li2O + 4NO2 + O2 ; 2Mg(NO3)2 ——→ 2MgO + 4NO2 + O2


1. Sodium Oxide, 

Na2O 2NaNO+ 10 Na ——→ 6Na 2O + N2 ; 2NaNO2 + 6 Na ——→ 4Na 2O + N2


NO2O + H2O ——→ 2NaOH

2. Sodium Peroxide, Na2O2 

It is formed by heating sodium in excess of air free form moisture  and carbon dioxide or in excess of pure oxygen.
S-Block Elements (Part -1) Chemistry Notes | EduRevProperties2Na 2O2 + 2H2O ——→ 4NaOH + O2; Na 2O2 + H2SO4 ——→ Na 2SO4 + H2O2

3. Sodium Hydroxide (Caustic Soda), NaOH 

It is one of the important chemicals and manufactured on a very large scale forming an important chemical industry. It is most conveniently manufactured by one of the following processes:

(a) Methods involving sodium carbonate as a starting material: Two methods are used. These are 

(i) Causticisation process (Gossage process)
Na2Co3 + Ca(Oh)⇆ CaCo3 + 2NaOH

(ii) Lowig’s process 

Na2CO3 + Fe2O3 ——→ 2NaFe2 + CO2; 2Na2FeO2 + H2O ——→ 2NaOH + Fe2O3

(b) Methods involving sodium chloride as starting material (electrolysis of brine)

At Cathod  2H2O + 2e ⇆ H+ 2OH; Na+ + OH-  ⇆ NaOH

At anode   S-Block Elements (Part -1) Chemistry Notes | EduRev


1. Strong alkali: NaOH ⇆ Na+ + OH-

(i) NaOH + HCl ——→ NaCl + H2O;

(ii) 2NaOH + CO2 ——→ Na2CO3 + H2O

(iii) Al2O+ 2NaOH ——→ 2NaAlO2 + H2O

2. Action on non-metals Cl2 + 2NaOH ——→ NaCl + NaClO + H2O

3Cl2 + 6NaOH ——→ 5NaCl + NaClO+ 3H2O

P4 + 3NaOH + 3HO ——→ 3NaH2PO2 + PH3

3. Action on Metals Zn + 2NaOH ——→ Na2ZnO2 + H2;

4. Action on Salts

(i) Ni(NO3)2 + 2NaOH ——→ Ni(OH)2 + 2NaNO3

(ii) Insoluble hydroxides which dissolve in excess of NaOH

ZnSO4 + 2NaOH ——→ Zn(OH)2 + Na2SO4

(iii) Unstable hydroxides 2AgNO3 + 2NaOH ——→ 2AgOH + 2NaNO3

2AgOH ——→ Ag2O + H2O


(a) Le-Blanc Process

(i) Conversion of NaCl into 

Na2CO4 NaCl + H2SO4 ——→ NaHSO4 + HCl

NaHSO4 + NaCl ——→ Na2SO4 + HCl

(b) Solvay ammonia soda process

NH3 + H2O + CO2 ——→ NH4HCO3

NaCl + NH4HCO3 ——→ NaHCO3 + NH4Cl

S-Block Elements (Part -1) Chemistry Notes | EduRev

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