The S-Block Elements

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Group 1 Elements Alkali Metals 

The elements of group 1 are known as alkali metals because their oxides and hydroxides are basic in nature. Aqueous solution of their oxides and hydroxides turn red litmus paper into blue. 


Elements of the First Group

Symbol 

Ions 

Color When Placed in a Flame 

Electronic Configuration 

Atomic and Ionic Radii 

Ionization Enthalpy 

Hydration Enthalpy (of ions)

Li 

Li+

Crimson red 

[He]2s1

Increases on moving from top to bottom in the group

Decreases on moving from top to bottom in the group

Decreases on moving from top to bottom in the group 

Na 

Na+ 

Yellow 

[Ne]3s1

K

K+

Violet 

[Ar]4s1

Rb 

Rb+ 

Red violet 

[Kr]5s1

Cs 

Cs+ 

Blue 

[Xe]6s1

Fr 

Fr+ 

-

[Rn]7s1



Elements of the First Group – Physical Properties 

Symbol 

Atomic Number 

Atomic Mass

(g mol-1)

Melting Point (K)

Boiling Point (K)

Density 

Ionic Radius 

Li 

3

6.94

454

Decreases on moving from top to bottom in the group

1615

Decreases on moving from top to bottom in the group

Increases on moving from top to bottom in the group (Exception – K-shows lower density)

Increases on moving from top to bottom 

Na 

11

22.99

371 

1156

K

19

39.10 

336

1032

Rb 

37

85.47 

312

961

Cs 

55

132.91 

302

944

Fr 

87

223

-

-



Elements of the First Group – Chemical Properties

Reactivity Towards Air 

Alkali metals react with air and form their oxides. These oxides react with moisture and form hydroxides. 

4Li + O2 🡪 2Li2O (monoxide)

2Na + O2 🡪 Na2O2 (peroxide)

M + O2 🡪 MO2 (Superoxide, where M = K, Rb, Cs)

Note – Alkali metals are highly reactive towards oxygen and water, so they are generally kept in kerosene. 

Reactivity Towards Water 

They easily react with water and form hydroxide and hydrogen. 

2M + 2H2O 🡪 2M+ +2OH- + H2

Note – Li is a small size alkali metal with very high hydration energy, so it reacts less vigorously with water.

Reactivity Towards Dihydrogen 

They react with dihydrogen and forms their respective hydrides. The reaction occurs at 673K temperature. 

2M + H2 🡪 2M+H- 

Note – Li reacts with H2 at 1073K temperature. 

Alkali metal hydrides show high melting points as they are ionic solids. 

Reactivity Towards Halogens 

They react vigorously and rapidly with halogens and form ionic halides. 

2M + X2 🡪 2MX

X = halogen 

Note – LiI is most covalent in nature because lithium ion has high polarization capability and is small in size. 

Reducing Nature 

All alkali metals are strong reducing agents. 

M + H2O 🡪 1/2H2 + MOH

(all alkali metals can reduce water to produce hydrogen gas)

Note – Li is the most powerful while Na is the least powerful reducing agent among alkali metals. 

Solutions in Liquid Ammonia 

All alkali metals can be dissolved in liquid ammonia. On standing, this solution slowly releases H2 and forms amide. 

M+(dissolved in ammonia) + e- + NH3 🡪 MNH2 + 1/2H2

Note - Alkali metals form deep blue colored paramagnetic solutions with ammonia. 



General Characteristics of the Compounds of the Alkali Metals 

Compounds

Characteristics of the Compounds

Oxides of Alkali Metals 

Alkali metals form monoxide, peroxide and superoxide. 

Stability of oxides increases as the size of the metal ion increases. 

Superoxides are either yellow or orange in color and paramagnetic in nature. 

In pure state oxides and peroxides are colorless. 

Hydroxides of Alkali Metals 

Hydroxides which are formed by the reaction of oxide and water are white colored crystalline solids. 

These hydroxides evolve a large amount of heat when dissolved in water. 

These hydroxides are the strongest bases. 

Halides of Alkali Metals 

These halides are colorless crystalline solids which possess high melting points. 

Melting and boiling points of alkali metal halides depend on the halogen ion. They follow the following trend – F- > Cl- > Br- > I-

These halides are soluble in water. 

Salts of Alkali Metals (with Oxo Acids)

When alkali metals react with oxo acids, they form salts. These salts are soluble in water and thermally stable. 

Stability of these salts increases as we move down the group. 

Lithium salts are exceptions. As lithium carbonate is thermally unstable and lithium hydrogen carbonate is not a crystalline solid salt. 


Anomalous Properties of Lithium 

Due to its exceptionally small size and high polarizing capacity, lithium shows different properties than other alkali metals. 


Difference Between Li and Other Alkali Metals 

S. No.

Lithium 

Other Alkali Metals 

1. 

It is harder. 

These are soft. 

2.

Its melting and boiling points are higher. 

Their melting and boiling points are lower than Li. 

3.

It is the least reactive alkali metal. 

All other alkali metals are more reactive than Li. 

4.

It is the strongest reducing agent among alkali metals. 

These are weaker reducing agents than LI.  

5.

Its chloride (LiCl) forms hydrate (LiCl.2H2O). lithium chloride is deliquescent while in its hydrate form, it’s a crystalline solid. 

Other alkali metals do not form hydrates. 

6.

Lithium hydrogen carbonate is found in solution form. 

Hydrogen carbonates of other alkali metals are found in solid state. 

7.

Lithium does not react with ethyne. 

Other alkali metals react with ethyne and form corresponding ethynide. 

8.

On decomposition, lithium nitrate gives lithium oxide. 

Reaction – 4LiNO3 🡪 2Li2O + 4NO2 + O2

Nitrates of other alkali metals on decomposition gives corresponding nitrite. 

Reaction – 2NaNO3 🡪 2NaNO2 + O2

9.

Lithium fluoride is much less soluble in water than other alkali fluorides. 

Fluorides of other alkali metals are soluble in water. 


Similarities Between Li and Mg– Lithium shows diagonal relationship with magnesium. 

  • Lithium is harder and lighter than other alkali metals. Magnesium is also harder and lighter than other elements of the second group. 

  • Both Li and Mg react slowly with water. 

  • Oxides and hydroxides of Li and Mg are less soluble in water and decompose on heating. 

  • Lithium nitride and magnesium nitride are formed by combination reaction with nitrogen. 

  • Their oxides do not produce superoxide. 

  • Lithium carbonate and magnesium carbonate produce oxides and carbon dioxide on decomposition. 

  • Chlorides of both are soluble in ethanol. 

  • Chlorides of both are deliquescent but their hydrates (LiCl.2H2O and MgCl2.8H2O) are crystalline solids. 


Some Important Compounds of Sodium 

Compound of Sodium 

Preparation 

Properties 

Uses 

Sodium carbonate
General name – washing soda 

Formula – Na2CO3.10H2O

It is prepared by the Solvay process. 

2NH3 + H2O + CO2 🡪 (NH4)2CO3

(NH4)2CO3 + H2O + CO2 🡪 2NH4HCO3

NH4HCO3 + NaCl 🡪 NH4Cl + NaHCO3

2 NaHCO3 🡪 Na2CO3 + CO2 + H2

  • White crystalline solid. 

  • Soluble in water.

  • On heating, it loses its water of crystallization

  • At higher temperatures, it changes into its anhydrous form. 

  • In its anhydrous form, it is found as white powder, which is called soda ash. 

  • It gets hydrolyzed by water and form an alkaline solution. 

  • As water softening reagent. 

  • In laundering and cleaning. 

  • In manufacturing of glass, soap, borax and caustic soda. 

  • It is one of the reagents used in textile, paint and paper industries. 

  • It is an important reagent for qualitative and quantitative analysis in laboratories. 

Sodium chloride

General name – Table salt, Common salt 

Formula – NaCl  

  • By crystallization of brine solution. 

  • By evaporation and crystallization of sea water.

  • Its melting point is 1081K. 

  • It is highly soluble in water. Its solubility in water is 36 g per 100 ml of water at 273 K. 

  • Change in temperature doesn’t affect its solubility in water up to large extent. 

  • Used in food items as common salt. 

  • Used in the production of sodium oxide, caustic soda, sodium carbonate. 

Sodium hydroxide
General name – Caustic soda

Formula – NaOH

  • By electrolysis method – It is prepared by electrolysis of brine solution (conc. solution of sodium chloride) in Castner – Kellner cell. 

At cathode – Na+ + e-Hg 🡪 Na amalgam 

At anode – Cl- 🡪 1/2Cl2 + e- 

2NaHg (sodium amalgam) + 2H2O 🡪 2NaOH + 2Hg + H2

  • It is white in color.

  • It is a translucent solid. 

  • Its melting point is 591 K.

  • It is soluble in water and gives an alkaline solution.

  • NaOH crystals are deliquescent.

  • Its aqueous solution reacts with carbon dioxide and gives sodium carbonate. 

  • In the manufacturing of soap.

  • Used in the production of paper, artificial silk and various chemicals. 

  • Used in petroleum refining. 

  • Used in the purification process of bauxite. 

  • Used in the cotton textile industry.

  • Used as laboratory reagent in various experiments.

  • Used for the preparation of pure fats and oils. 

Sodium hydrogencarbonate
General name – Baking soda 

Formula – NaHCO3

  • By saturation of the solution of sodium carbonate with carbon dioxide.

Na2CO3 + H2O + CO2 🡪 2NaHCO3

  • It is a white crystalline solid. 

  • It is an odorless substance. 

  • Its melting point is 323 K. 

  • It is used in baking as it decomposes on heating and generates carbon dioxide which makes the baked dishes spongy. 

  • It is used as an antiseptic for skin infections. 

  • It is used in fire extinguishers.  


Biological Importance of Sodium and Potassium 

Sodium and potassium elements of the group 1 are of biological importance. They are important minerals of our body. Their biological importance can be described by the following points –

  • A human body contains approximately 0.12% of Na and 0.25% of K. 

  • Sodium ions are found in blood plasma and interstitial fluid.

  • Sodium ions help in transporting the signals from nerve cells, regulating the water flow across cell membranes, transport of sugars and amino acids to cells.

  • Sodium ions also help in activation of various enzymes.  

  • Potassium ions also help in the transmission of nerve signals and transportation of essential compounds to cells. 

  • Potassium ions are the most abundant positive ions within cell fluids. 

  • Potassium ions activate many enzymes which help in production of energy by oxidation of glucose. 

  • Sodium – potassium pump which operates across the cell membrane is based on the sodium and potassium ions. 

  • Sodium and potassium ions differ in their concentrations in the cell fluid and also differ quantitatively in their ability to penetrate the cell membrane. 


The elements of group 2 are known as alkaline earth metals because their oxides and hydroxides are basic in nature and these metals are found in earth or earth’s crust. Aqueous solution of their oxides and hydroxides turn red litmus paper into blue. 

Group 2 Elements: Alkaline Earth Metals 

Elements of the Second Group

Atomic Number (Z)

Symbol 

Ions 

Color When Placed in a Flame 

Electronic Configuration 

Atomic and Ionic Radii 

Ionization Enthalpy 

Hydration Enthalpy (of ions)

4

Be 

Be+2

White 

[He]2s2

Increases on moving from top to bottom in the group

Decreases on moving from top to bottom in the group due to increase in size.

Decreases on moving from top to bottom in the group. 

12

Mg 

Mg+2 

White 

[Ne]3s2

20

Ca 

Ca+2

Brick red 

[Ar]4s2

38

Sr 

Sr+2 

Crimson 

[Kr]5s2

56

Ba 

Ba+2 

Apple green 

[Xe]6s2

88

Ra 

Ra+2 

Crimson red 

[Rn]7s2



Elements of the Second Group – Physical Properties 

Symbol 

Atomic Number 

Atomic Mass

 (g mol-1)

Melting Point (K)

Boiling Point (K)

Density 

Ionic Radius 

Be 

4

9.01

1560 

Decreases on moving from top to bottom in the group (Mg and Ra have exceptionally low m.p.)

2745 

Decreases on moving from top to bottom in the group (Mg has exceptionally low and Ba has exceptionally high b.p.)

Decreases on moving till calcium and then increases till radium. 

Increases on moving from top to bottom 

Mg 

12

24.31 

924 

1363

Ca 

20

40.08

1124 

1767

Sr 

38

87.62

1062

1655

Ba 

56

137.33

1002

2078 

Ra 

88

226.03

973

1973



Elements of the Second Group – Chemical Properties

Reactivity Towards Air 

Alkaline earth metals react with air and form their oxides. These oxides react with moisture and form hydroxides. 

On reacting with air, they form nitrides as well. 

2Mg + O2 🡪 2MgO 

2Be + O2 🡪 BeO 

2M + O2 🡪 2MO 

Note – Alkaline earth metals are less reactive than alkali metals, although reactivity of alkaline earth metals increases on moving down the group. 

Reactivity Towards Water 

They easily react with water and form hydroxide and hydrogen. 

Ca + 2H2O 🡪 Ca(OH)2 + H2

M + 2H2O 🡪 M(OH)2 + H2

(where M = Ca, Sr, Ba)

Note – Be and Mg are less reactive towards water although their oxides readily react with water. 

Reactivity Towards Dihydrogen 

They react with dihydrogen and form their respective hydrides. Their hydrides are unstable in water. 

M + H2 🡪 MH2 

Note – Their hydrides are saline in nature.  


Reactivity Towards Halogens 

All alkaline metals react with halogens and form ionic halides. The reaction takes place at high temperature. 

M + X2 🡪 MX2

X = F, Cl, Br, I


Note – Calcium chloride is hygroscopic in nature and on exposure to air, it absorbs water and forms a solution. 

Reducing Nature 

All alkaline earth metals are strong reducing agents but weaker than alkali metals. 

M + H2O 🡪 H2 + M(OH)2

(all alkali metals can reduce water to produce hydrogen gas, except Be)

Note – Be has least reducing nature among alkaline earth metals. 

Solutions in Liquid Ammonia 

All alkaline earth metals can be dissolved in liquid ammonia. They form a deep blue - black colored solution. 

Ca + 2NH3 🡪 Ca(NH2)2 + H2

Note - Alkali metals form deep blue - black colored solution with ammonia. 

Reaction with Acids 

Alkaline earth metals react with acids and releases hydrogen gas. 

M + 2HCl 🡪 MCl2 + H2

Note – M can be any alkaline earth metal. 

Reaction with Nitrogen 

All alkaline earth metals do not react with nitrogen directly. Only Be and Mg react with N2 directly. 

3Be + N2 🡪 Be3N2

3Mg + N2 🡪 Mg3N2

Note – Be and Mg on reaction with air can form nitrides directly if enough amount of nitrogen is present in the air in the area.   

Reaction with Alkyl Halide 

Mg reacts with alkyl halide through an insertion reaction or combination reaction. 

RX + Mg 🡪 RMgX 

Note - RMgX is known as Grignard reagent. 



General Characteristics of Compounds of the Alkaline Earth Metals

Types of Compounds 

Characteristics of the Compounds of Alkaline Earth Metals 

Oxides of Alkaline Earth Metals 

The alkaline earth metals react with oxygen and form corresponding oxides. 

They form monoxide. 

BeO is being an exception is covalent in nature. It is an amphoteric oxide. While oxides of other alkaline earth metals are ionic in nature.  

2M + O2 🡪 2MO

  • These oxides possess high enthalpy of formation.

  • They are thermally stable. 

Hydroxides of Alkaline Earth Metals 

Except BeO, all other alkaline earth metal oxides are basic in nature and form their respective hydroxides with water. 

Solubility, thermal stability and basic character of hydroxides of alkaline earth metals increases from Mg to Ba due to increase in atomic size. 

MO + H2O 🡪 M(OH)2

  • Beryllium hydroxide is amphoteric in nature which means it reacts with acid and base both. 

  • Hydroxides of alkaline earth metals are less basic and stable than hydroxides of alkali metals. 

Halides of Alkaline Earth Metals 

Halides of all alkaline earth metals are ionic in nature, except BeX2

The tendency of forming halide hydrates decreases on moving down the group (from Mg to Ba)

Fluorides of alkaline earth metals are relatively less soluble than chlorides of alkaline earth metals. 

Salts of Oxoacids – Carbonates 

  • Alkaline earth metals form salts with oxoacids. 

  • Carbonates of alkaline earth metals are water insoluble. 

These can be precipitated by addition of ammonium carbonate and sodium carbonate. 

  • Solubility and thermal stability of carbonates of alkaline earth metals increases on moving down the group due increase in the size of metal ions. 

  • Beryllium carbonate is unstable and should be kept in the atmosphere of CO2

Salts of Oxoacids – Sulphates 

  • Sulphates of alkaline earth metals are white crystalline solids and thermally stable. 

Solubility of sulphates of alkaline earth metals decreases as we move down the group. Although BeSO4 and MgSO4 show almost the same solubility in water. 

  • Beryllium and magnesium sulphates are readily soluble in water due to their high hydration enthalpies. 

Salts of Oxoacids – Nitrates 

  • Nitrates of alkaline earth metals are formed by dissolution of the carbonates in dil. HNO3

Nitrates of all alkaline earth metals decompose on heating and give their respective oxides. 

2M(NO3)2 🡪 2MO + 4NO2 + O2

(Where M = Be, Mg, Ca, Sr or Ba)

  • Crystalline form of magnesium nitrate has six crystallization of water molecules. 


Anomalous Behavior of Beryllium 

Due to its exceptionally small size and high ionization enthalpies, beryllium shows different properties than other alkaline earth metals or 2nd group elements.  


Difference Between Be and Other Alkaline Earth Metals 

S. No.

Be (Beryllium)

Other Alkaline Earth Metals 

1.

Compounds of Be are covalent in nature. 

Compounds of other alkaline earth metals are ionic in nature. 

2.

Compounds of Be, easily get hydrolyzed. 

Compounds of other alkaline do not get hydrolyzed easily. 

3.

Beryllium carbonate is unstable in nature. 

Carbonates of other alkali earth metals are stable in nature. 

4.

Oxide of beryllium is amphoteric in nature. 

Oxides of all other alkaline earth metals are basic in nature. 

5.

Be does not possess coordination number more than four. 

Other alkaline earth metals exhibit coordination number up to six. 

6.

Beryllium sulphate is readily soluble in water. 

Sulphates of other alkaline earth metals (Except magnesium) possess less solubility in water than BeSO4

7.

Beryllium hydroxide is amphoteric in nature. 

Hydroxides of all other alkaline earth metals are basic in nature.


Similarities Between Be and Al – Beryllium shows diagonal relationship with aluminum.  Some similarities between the are listed below –

  • Aluminium and beryllium both do not attack acids easily. 

  • Beryllium hydroxide and aluminium hydroxide both give beryllium ion and aluminium ion respectively on dissolving in excess of aqueous solution of base. 

  • In vapor phase chlorides of both Al and Be show chloride bridge structure. 

  • Chlorides of both are strong Lewis acids and soluble in organic solvents. 

  • Ions of both form complexes such as BeF42-, AlF63- etc. 


Some Important Compounds of Calcium 

Compound of Calcium 

Preparation 

Properties 

Uses 

Calcium oxide 

Common name – Quicklime 

Formula – CaO 

By heating calcium carbonate in a rotary kiln at 1070 – 1200 K. 

CaCO3 Δ↔️ CaO + CO2

  • It is white amorphous solid. 

  • It melts at 2870 K. 

  • It can absorb moisture from air. 

CaO + H2O 🡪 Ca(OH)2

  • It can absorb CO2 from air. 

CaO + CO2 🡪 CaCO3

  • In white washing. 

  • In manufacturing of cement. 

  • In production of sodium carbonate. 

  • Used in the purification of sugar. 

  • In the manufacturing of dye stuffs. 

Calcium hydroxide 

Common name – Slaked lime 

Formula – Ca(OH)2

It is prepared by reaction of quick lime with water. 

CaO + H2O 🡪 Ca(OH)2

  • It is also a white amorphous powder. 

  • It is slightly soluble in water. 

  • Lime water – Aqueous solution of Ca(OH)2.

  • Milk of lime – Suspension of calcium hydroxide in water. 

  • It reacts with carbon dioxide and water and gives calcium hydrogen carbonate. 

  • It reacts with chlorine and forms bleaching powder. 

  • Used in preparation of mortar and bleaching powder. 

  • It is used in whitewash. 

  • Used in glass making, purification of sugar and tanning industry. 

Calcium carbonate 

Common name – Limestone 

Formula – CaCO3


  • By passing carbon dioxide through calcium hydroxide. 

Ca(OH)2 + CO2 🡪 CaCO3 + H2

  • By reaction of calcium chloride and sodium carbonate. 

CaCl2 + Na2CO3 🡪 CaCO3 + 2NaCl

  • It is a white fluffy powder. 

  • It is insoluble in water. 

  • At high temperatures, it decomposes to give calcium oxide and carbon dioxide. 

CaCO3 🡪 CaO + CO2

  • Used as decorative and building material in buildings. 

  • Used as flux in extraction of metals. 

  • Used as an antacid to relief acidity. 

  • Used in toothpaste, chewing gum, cosmetic materials.  

Calcium sulphate 

Common name – Plaster of Paris

Formula – CaSO4.1/2H2O  

By heating gypsum at 393 K. 

CaSO4.2H2O 393K 🡪 CaSO4.1/2H2O + 3/2H2O

  • It is a hemihydrate of calcium sulphate. 

  • On heating, it loses its water of crystallization and forms its anhydrous form which is called ‘dead burnt plaster’.

  • Used in treatment of bone fractures. 

  • Used in buildings. 

  • Used in dental treatment. 

  • Used in ornamental work. 


Biological Importance of Magnesium and Calcium 

Magnesium and calcium are the elements of the group 1 which has biological importance. They are important minerals of our body. Their biological importance can be described by the following points –

  • Calcium is necessary for our healthy bones. Almost 1200g of calcium is found in the body of an adult. 

  • Magnesium is important for various enzymes involved in the utilization of energy (ATP). Almost 25g of Mg is found in the body of an adult human being. 

  • Chlorophyll which is found in the leaves contains Mg. 

  • Ca is necessary for proper growth of our body. 

  • Ca is important for healthy teeth as well.

  • Calcium plays a vital role in blood coagulation, neuromuscular function and building of cell walls in plant cells. 

This ends our coverage on the summary of the unit “The s-block elements”. We hope you enjoyed learning and were able to grasp the concepts. You can get separate articles as well on various subtopics of this unit such as alkali metals, elements of group 2 etc. on Vedantu website. We hope after reading this article you will be able to solve problems based on the topic. We have already provided detailed study notes or revision notes for this unit, which you can easily download by registering yourself on Vedantu website. Here in this article we have discussed the unit in a summarized way with the emphasis on important topics of the unit.  If you are looking for solutions of NCERT Textbook problems based on this topic, then log on to Vedantu website or download Vedantu Learning App. By doing so, you will be able to access free PDFs of NCERT Solutions as well as Revision notes, Mock Tests and much more.