THE S-BLOCK ELEMENTS
The elements in which the last electron enters in the outermost s-orbital are called s-block elements.
They include elements of group I and II. Group I elements are Lithium (Li), Sodium (Na), Potassium (K),
Rubidium (Rb), Caesium (Cs) and Francium (Fr). They are collectively called alkali metals because they dissolve
in water to form strong alkalies.
Group II elements are Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba) and
Radium (Ra). They are collectively called alkaline earth metals (except Be) because their oxides and hydroxides
are found in earth crust and alkaline in nature.
The general electronic configuration of s-block elements is [noble gas] ns1 for alkali metals and [noble
gas] ns2 for alkaline earth metals.
GROUP I ELEMENTS [ALKALI METALS]
General characteristics
1. Valence electronic configuration: ns1
2. Atomic and ionic radii: Increases from top to bottom.
3. Ionization enthalpy: decrease from Li to Cs.
4. Hydration enthalpy: decreases with increase in ionic size. Li+ has maximum hydration enthalpy.
5. Flame colouration: Alkali metals and their salts give characteristic colour to non-luminous flame. This is
because the heat from the flame excites the outer most orbital electron to a higher energy level. When this
electron comes back to the ground level, they emit the radiation in the visible region. For example, Li gives
crimson red, sodium gives yellow, potassium gives violet, Rubidium gives red violet and Ceasium gives blue
colour to the flame. So alkali metals can be detected by flame test.
Chemical Properties
1. Reaction with air: They react with air to form oxides, peroxides and super oxides. Li forms only
monoxide, sodium forms monoxide and peroxide and other alkali metals form monoxide, peroxide and
super oxide.
4Li + O2 → 2Li2O
2Na + O2 → Na2O2
M + O2 → MO2 (where M = K, Rb, Cs)
2. Reaction with water: The alkali metals react with water to liberate hydrogen.
2M + H2O → 2MOH + H2
But the reactivity of Li with water is less vigorous due to its small size and very high hydration
enthalpy.
3. Reaction with Hydrogen:
2M + H2 → 2MH (metal hydride)
4. Reaction with halogen:
2M + X2 → 2MX
5. Reducing nature: They are strong reducing agents. Li being the most powerful reducing agent and
sodium the least. Due to the smallest atomic radius, Li has the highest hydration enthalpy. So it has high
reducing power.
6. Solution in liquid ammonia:
The alkali metals dissolve in liquid ammonia to give deep blue solutions which are good conductors.
M + (x+y)NH3 → [M(NH3)x]+ + e[(NH3)y]-
The blue colour of the solution is due to the ammoniated electron, which absorbs energy in the visible
region and gives blue colour to the solution. The solution is paramagnetic and on standing slowly liberates
hydrogen resulting in the formation of amide (MNH2).
[M(NH3)x]+ → MNH2 + ½ H2
In concentrated solution, the blue colour changes to bronze colour and become diamagnetic.
The s-block elements
, Anomalous Properties of Lithium
Due to its small size and high polarizing power, Lithium shows some properties different from that of other
alkali metals. Some of these are:
1. Li is much harder and has high melting point and boiling point.
2. Li is the least reactive but the strongest reducing agent among all the alkali metals.
3. It forms only monoxide with oxygen.
4. LiCl is deliquescent and crystallizes as a hydrate (LiCl.2H2O). But the other alkali metal chlorides do not
form hydrates.
5. Lithium bicarbonate (LiHCO3) is stable only in solution.
6. Lithium nitrate on heating gives Li2O, while other alkali metal nitrates decompose to form nitrite.
4 LiNO3 → 2Li2O + 4NO2 +O2
2NaNO3 → 2NaNO2 + O2
Diagonal relationship
The similarity in properties shown by diagonally placed elements of second and third periods in
modern periodic table is called diagonal relationship.
Diagonal relationship between Li and Mg
Li shows the following similarities in properties with Be of the second group.
1. Both Li and Be are harder but lighter than other elements of the respective group.
2. Both react slowly with water. Their oxides and hydroxides are much less soluble and their hydroxides
decompose on heating.
3. They do not form superoxides.
4. Their carbonates decompose easily on heating to form oxides and CO2.
5. Their chlorides are soluble in ethanol and are deliquescent.
6. Their bicarbonates are stable only in solution.
Some important compounds of Sodium
1. Sodium Carbonate [Na2CO3.10H2O] (Washing Soda)
Preparation: Solvay Process (Ammonia-Soda Process)
In this process, CO2 is passed through a concentrated solution of NaCl saturated with ammonia.
Ammonium carbonate first formed then converted to ammonium bicarbonate and finally reacts with NaCl to
form NaHCO3.
2NH3 + H2O + CO2 → (NH4)2CO3
(NH4)2CO3 + H2O + CO2 → 2NH4HCO3
NH4HCO3 + NaCl → NH4Cl + NaHCO3
Sodium bicarbonate crystals are separated and heated to get sodium carbonate.
2NaHCO3 → Na2CO3 + CO2 + H2O
In this process, NH3 is recovered when the solution containing NH4Cl is treated with Ca(OH)2.
2NH4Cl + Ca(OH)2 → 2NH3 + CaCl2 + 2H2O
Note: Solvay process cannot be used for the preparation of K2CO3 because potassium bicarbonate (KHCO3) is
so much soluble in water that it does not get precipitated by the addition of NH4HCO3 to a saturated solution
of KCl.
Properties:
Action of Heat: On heating, the decahydrate ( Na2CO3.10H2O) loses its water of crystallisation to form a
monohydrate. Above 373K the monohydrate becomes completely anhydrous and changes to a white powder
called “soda ash”.
Na2CO3.10H2O → Na2CO3.H2O + 9H2O
Na2CO3.H2O → Na2CO3 + H2O
The s-block elements
The elements in which the last electron enters in the outermost s-orbital are called s-block elements.
They include elements of group I and II. Group I elements are Lithium (Li), Sodium (Na), Potassium (K),
Rubidium (Rb), Caesium (Cs) and Francium (Fr). They are collectively called alkali metals because they dissolve
in water to form strong alkalies.
Group II elements are Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba) and
Radium (Ra). They are collectively called alkaline earth metals (except Be) because their oxides and hydroxides
are found in earth crust and alkaline in nature.
The general electronic configuration of s-block elements is [noble gas] ns1 for alkali metals and [noble
gas] ns2 for alkaline earth metals.
GROUP I ELEMENTS [ALKALI METALS]
General characteristics
1. Valence electronic configuration: ns1
2. Atomic and ionic radii: Increases from top to bottom.
3. Ionization enthalpy: decrease from Li to Cs.
4. Hydration enthalpy: decreases with increase in ionic size. Li+ has maximum hydration enthalpy.
5. Flame colouration: Alkali metals and their salts give characteristic colour to non-luminous flame. This is
because the heat from the flame excites the outer most orbital electron to a higher energy level. When this
electron comes back to the ground level, they emit the radiation in the visible region. For example, Li gives
crimson red, sodium gives yellow, potassium gives violet, Rubidium gives red violet and Ceasium gives blue
colour to the flame. So alkali metals can be detected by flame test.
Chemical Properties
1. Reaction with air: They react with air to form oxides, peroxides and super oxides. Li forms only
monoxide, sodium forms monoxide and peroxide and other alkali metals form monoxide, peroxide and
super oxide.
4Li + O2 → 2Li2O
2Na + O2 → Na2O2
M + O2 → MO2 (where M = K, Rb, Cs)
2. Reaction with water: The alkali metals react with water to liberate hydrogen.
2M + H2O → 2MOH + H2
But the reactivity of Li with water is less vigorous due to its small size and very high hydration
enthalpy.
3. Reaction with Hydrogen:
2M + H2 → 2MH (metal hydride)
4. Reaction with halogen:
2M + X2 → 2MX
5. Reducing nature: They are strong reducing agents. Li being the most powerful reducing agent and
sodium the least. Due to the smallest atomic radius, Li has the highest hydration enthalpy. So it has high
reducing power.
6. Solution in liquid ammonia:
The alkali metals dissolve in liquid ammonia to give deep blue solutions which are good conductors.
M + (x+y)NH3 → [M(NH3)x]+ + e[(NH3)y]-
The blue colour of the solution is due to the ammoniated electron, which absorbs energy in the visible
region and gives blue colour to the solution. The solution is paramagnetic and on standing slowly liberates
hydrogen resulting in the formation of amide (MNH2).
[M(NH3)x]+ → MNH2 + ½ H2
In concentrated solution, the blue colour changes to bronze colour and become diamagnetic.
The s-block elements
, Anomalous Properties of Lithium
Due to its small size and high polarizing power, Lithium shows some properties different from that of other
alkali metals. Some of these are:
1. Li is much harder and has high melting point and boiling point.
2. Li is the least reactive but the strongest reducing agent among all the alkali metals.
3. It forms only monoxide with oxygen.
4. LiCl is deliquescent and crystallizes as a hydrate (LiCl.2H2O). But the other alkali metal chlorides do not
form hydrates.
5. Lithium bicarbonate (LiHCO3) is stable only in solution.
6. Lithium nitrate on heating gives Li2O, while other alkali metal nitrates decompose to form nitrite.
4 LiNO3 → 2Li2O + 4NO2 +O2
2NaNO3 → 2NaNO2 + O2
Diagonal relationship
The similarity in properties shown by diagonally placed elements of second and third periods in
modern periodic table is called diagonal relationship.
Diagonal relationship between Li and Mg
Li shows the following similarities in properties with Be of the second group.
1. Both Li and Be are harder but lighter than other elements of the respective group.
2. Both react slowly with water. Their oxides and hydroxides are much less soluble and their hydroxides
decompose on heating.
3. They do not form superoxides.
4. Their carbonates decompose easily on heating to form oxides and CO2.
5. Their chlorides are soluble in ethanol and are deliquescent.
6. Their bicarbonates are stable only in solution.
Some important compounds of Sodium
1. Sodium Carbonate [Na2CO3.10H2O] (Washing Soda)
Preparation: Solvay Process (Ammonia-Soda Process)
In this process, CO2 is passed through a concentrated solution of NaCl saturated with ammonia.
Ammonium carbonate first formed then converted to ammonium bicarbonate and finally reacts with NaCl to
form NaHCO3.
2NH3 + H2O + CO2 → (NH4)2CO3
(NH4)2CO3 + H2O + CO2 → 2NH4HCO3
NH4HCO3 + NaCl → NH4Cl + NaHCO3
Sodium bicarbonate crystals are separated and heated to get sodium carbonate.
2NaHCO3 → Na2CO3 + CO2 + H2O
In this process, NH3 is recovered when the solution containing NH4Cl is treated with Ca(OH)2.
2NH4Cl + Ca(OH)2 → 2NH3 + CaCl2 + 2H2O
Note: Solvay process cannot be used for the preparation of K2CO3 because potassium bicarbonate (KHCO3) is
so much soluble in water that it does not get precipitated by the addition of NH4HCO3 to a saturated solution
of KCl.
Properties:
Action of Heat: On heating, the decahydrate ( Na2CO3.10H2O) loses its water of crystallisation to form a
monohydrate. Above 373K the monohydrate becomes completely anhydrous and changes to a white powder
called “soda ash”.
Na2CO3.10H2O → Na2CO3.H2O + 9H2O
Na2CO3.H2O → Na2CO3 + H2O
The s-block elements
