The S Block Elements Class 11 Notes CBSE Chemistry Chapter 10

The S-Block Elements Class 11 Notes for Chemistry Chapter 10

The s-Block elements are the Group 1 and Group 2 elements of the periodic table which are called Alkali Metals and Alkali Earth Metals respectively. These are characterised by one s-electron for Alkali Metals and two s- electrons for Alkali Earth Metals. 

Extramarks, one of the leading  e-learning platforms, is known for providing the best educational materials for students of all classes. Extramarks Revision Notes for Class 11 Chemistry Chapter 10 provides comprehensive notes of each element’s properties and reactions, which helps students prepare for their exams.

Access Class 11 Chemistry Chapter 10 Notes – The S Block Elements

Introduction

Students will be able to explain the fundamental properties of alkali metals and their compounds in this chapter, as well as the fundamental properties of alkaline earth metals and their compounds. They will also be able to describe how sodium and calcium compounds, which are crucial for the industry, are made and used, as well as their properties and uses, and understand the biological significance of sodium, potassium, magnesium, and calcium.

The s-Block elements are categorised into two groups as given below.

Alkali Metals (Group 1)

These are highly reactive and have ns¹ configuration.

Physical Properties

  • Atomic Size: The atomic size increases across the group. 
  • Oxidation State: Group 1 elements have a +1 Oxidation State.
  • Density: They have low density due to their large size. Hence density rises from Li to Cs.
  • Nature of Bonds: They combine with other elements when electronegativity is low to form ionic bonds.
  • Ionisation Energy: They have the lowest Ionisation Energy than any other group because outer electrons are weakly bound by the nucleus due to large atoms. IE decreases as we go through the group.
  • Flame Test: If an atom is excited by strong heat, the electrons in its valence shell can be promoted to higher orbitals. As they fall back down in their orbits in normal conditions, they emit energy as light, which gives flame colour. 
Elements Atomic Number Electronic Configuration Colour
Lithium 3 [He] 2s¹ Red
Sodium 11 [Ne] 3s¹ Golden yellow
Potassium 19 [Ar] 4s¹ Violet
Rubidium 37 [Kr] 5s¹ Red violet
Cesium 55 [Xe] 6s¹ Blue
Francium 87 [Rn] 7s¹

Standard Oxidation Potential

The standard electrode potential is defined as the concentration of metal ions being equal to one. It is the tendency of a species to be oxidised at standard conditions. Lithium has the largest Ionisation potential and highest electrode potential.

  • Hydration of Ions: the degree of hydration is proportional to the size of ions. Hence from Li to Cs degree of hydration decreases and electrical conductivity increases.
  • Lattice Energy: lattice energy decreases as we go down the group.
  • Solubility in Liquid Ammonia: Alkali Metals dissolve readily and give a blue colour due to the presence of ammoniated electrons. The colour vanishes slowly if left standing due to the formation of metal amide. Dilute solutions are paramagnetic. 
  • Electronegativity Values: From lithium to caesium, the electronegativity levels are low.
  • Reactivity: Alkali metals’ reactivity rises steadily in the following order: Li<Na<K<Rb<Cs

Colourless and Diamagnetic Ions

An ion’s colour is determined by its number of unpaired electrons. If an anion containing unpaired electrons is stimulated and returned to the ground state, it shows colour. Unpaired electrons have magnetic properties, while paired electrons do not and are called Diamagnetic ions. 

Chemical Properties

Due to their large size and low ionization enthalpy, alkali metals are highly reactive. As one moves down the group, the reactivity of these metals rises.

Common reactions-

  • With Air: Very reactive and tarnish quickly with air exposure. They burn in oxygen so that  lithium forms monoxide, sodium forms peroxide, and other metals form superoxides.
  • With water: They react with water to form hydroxide and dihydrogen.
  • With Dihydrogen: reaction with dihydrogen occurs at about 673K (lithium at 1073 K) to form hydrides.
  • With Halogens: react strongly with halogens to form ionic halides, M+X
  • Solution in liquid ammonia: they dissolve in liquid ammonia to give deep blue solutions which is are conducting in nature.

M+(x+y)NH3→[M(NH3)x]++[e(NH3)y]

Compounds of Alkali Metal

  • Hydroxides: Sodium hydroxide or Caustic Soda and potassium hydroxide or Caustic Potash are known for their corrosive nature. The solubility of hydroxides rises through the group.

Oxides, Peroxides and Superoxides

  • Monoxides: ionic monoxides are very basic oxides that form strong bases in the presence of water.
  • Peroxides: Na2O2  is a powerful oxidant as it reacts with carbon dioxides present in the air.
  • Superoxides: Superoxides are stronger oxidising agents than peroxides. 

The stability of oxides-

Normal oxide > peroxide > superoxide

Carbonates and Bicarbonates

Due to their electropositive nature, alkali metal carbonates and bicarbonates are exceptionally heat stable. Li2Co3, for instance, is easily broken down by heat.

Explanation: 

  • The Electron cloud of Oxygen from CO23 is disrupted due to Lithium’s small size and high polarization. It weakens the carbon-hydrogen bond. 
  • The Lattice Energy increases when a larger carbonate ion is replaced by a smaller carbonate ion. 

Halides:

All metals in group 1 create MX halides. Out of all these elements, Lithium is more likely to produce hydrated salts as it is the smallest.

  • Properties: Alkali metal halides are excellent ionic compounds as shown by the given characteristics.
    • Alkali Halides are easily soluble in water(except lithium fluoride)
    • MP and BP are high. They drop in the order of Fluoride > chloride > bromide > iodide.

                    The MP of lithium halides is lower than that of sodium halides. Going through the group          from sodium to caesium, the MP of halides decline as lattice energy decreases. 

  • Solubility of halides of Alkali Metals: Alkali Metals have a wide range of solubilities. When the lattice energy of the crystals decreases, solubility in water increases steadily from sodium chloride to caesium chloride.
  • Good conductors of electricity while fused. 
  • Made up of ionic crystals except for Lithium halides which have a partially covalent character.
  • Lattice energy and polarising power are responsible for the structure and solubility of alkali metal halides.

Anomalous Behaviour of Lithium

Lithium atoms and their ions exhibit peculiar behaviour, such as having a high charge density, as a result of their incredibly small size. The strongest polarising power, then, belongs to lithium ions, which are comparable to magnesium ions in terms of their characteristics. 

     Lithium ions have a strong inclination for solvation and the creation of covalent bonds. Its properties are given below:

  • Comparatively difficult to work with
  • High MP and BP
  • Least reactive unlike other elements in Group 1 since it is not influenced by air, slowly decomposes water, rarely reacts with bromine and forms only monoxide(Li2O)when burned in oxygen.
  • Forms nitride(Li3N)when in contact with nitrogen.
  • Comparatively less electropositive, hence several of its compounds are unstable. 
  • When heated, lithium nitrate produces nitrogen dioxide and oxygen, leaving lithium oxide while other elements produce only oxygen, leaving nitrites. 
  • Majority of lithium salts(hydroxide, carbonate, oxalate, etc) are water-insoluble while others are soluble.
  • Lithium halides and lithium alkyls are soluble in organic solvents while others aren’t.
  • Lithium chloride undergoes hydrolysis in water, unlike other group-1 elements. 
  • Partially covalent lithium halides exist in nature. Due to its high polarising power, lithium compounds have a smaller dipole moment.
  • Alkali metals have more hydrated ions and compounds than other alkali metals.

Alkaline Earth Metals

Introduction:

The metals barium, beryllium, calcium, magnesium, radium, and strontium, are all found in Group II of the periodic table. Because of its radioactivity, radium has chemical characteristics in common with alkaline earth metals.

Physical properties

  • Atomic Size: It increases as we go down the group.
  • Oxidation State: They have a +2 Oxidation State
  • Density: Density is high because group 2 components are smaller and have a higher density. As we descend the group, density increases.
  • Nature of Bonds: Elements for ionic bonds except for beryllium which forms covalent bonds.
  • Hydration Energy: Hydration energies of group 2 ions are 4-5 times higher than those of group 1 ions. Hydration enthalpy falls as the size of ions increases.
  • Lattice Energy: It decreases as we go down the group.
  • Ionisation Energy: Electrons are more closely bonded as the elements are smaller, hence more Ionisation Energy is required, about 4 times more than that of group 1 elements. 
  • Flame Test:

Calcium- Brick red

Strontium- Crimson red

Barium- Grassy green

Radium- Crimson

Due to the smaller size and more firmly linked atoms of beryllium and magnesium, these elements are not ignited by flame and do not emit any colour.

Standard Oxidation Potential

Elements Oxidation Reaction Standard Oxidation Potential (volt)
Be Be→ Be2+ + 2e 1.85
Mg Mg → Mg2+ + 2e 2.37
Ca Ca→ Ca2+ + 2e 2.87
Sr Sr→ Sr²+ + 2e 2.89
Ba Ba →Ba2+ + 2e 2.90

Solubility in Liquid Ammonia

The metals dissolve in liquid ammonia and their dilute solutions are blue in colour. Amides are formed and Hydrogen gas is released when it decomposes.

  • Electronegativity  Values: It is low but higher than group 1. Value decreases as we move down the group.
  • Colourless and Diamagnetism: Elements that have unpaired electrons generate M²+ ions which are Diamagnetic and Colourless.
  • Melting and Boiling Point: MP lowers as the group progresses.
  • Metallic Properties: Metallic sheen with electrical and thermal conductivity. 

Sodium Peroxides

  • Preparation: These are formed by heating the element in excess of oxygen at 300°C without CO2.
  • Properties: 
    • Pale yellow solid that becomes white in the air. 
    • It creates oxygen at normal temperature, hydrogen peroxide in ice-cold mineral acids, and H2O2 in cold water.
    • Combines with carbon dioxide to produce sodium carbonate and oxygen
    • It is an oxidising agent and oxidised charcoal, CO, NH3, SO2.
    • Contains peroxide ions.
  • Uses
    • For preparing H2O2, O2
    • Oxygenating the air in submarines
    • Oxidising agent in the laboratory.

Oxides of Potassium

Oxides of Potassium Colour 
K2O White
K2O2 White
K2O3 Red
KO2 Bright yellow
KO3 Reddish brown needles

Magnesium Oxide:

Properties

  • Present as a white powder.
  • MP over 2850°C. Hence, it is used for manufacturing refractory bricks for furnaces.
  • Slightly soluble in water.

Magnesium Peroxide and Calcium Peroxide:

These are obtained by passing H2O2 in a suspension of Mg(OH)2 and Ca(OH)2

  • Uses: Used as an antimicrobial and whitening agent in toothpaste. 

Potassium Hydroxide

  • Preparation: Electrolysis is used to prepare the potassium chloride aqueous solution.
  • Properties: 
    • Stronger base compared to sodium hydroxide.
    • Greater water solubility than sodium hydroxide.
    • KOH is highly soluble in alcohol.

Magnesium Hydroxide

It can be found as the mineral brucite in its natural form.

  • Preparation: Mix a magnesium sulphate or chloride solution with a caustic soda solution.
  • Properties: 
    • Can only be dried at temperatures up to 100 degrees Celsius, otherwise it breaks down to its oxide.
    • Mildly soluble in water, so it is alkalinizing. 
    • It dissolves in an NH4Cl solution.

Calcium Hydroxide

  • Preparation: It is prepared by spraying water on quicklime.
  • Properties:
    • Solubility in water is minuscule.
    • Solubility rises with the temperature of the water.
    • Carbon dioxide is readily absorbed by calcium hydroxide.

Potassium Carbonate

It can only be made by the Leblanc method since potassium carbonate is water soluble.

  • Properties: resembles Na2CO3 as its melting point is 900∘C.
  • Uses: used in glass manufacturing.

Calcium Carbonate

Natural calcite can be found in materials including marble, limestone, chalk, and calcite.

  • Properties: 
    • It dissociates at temperatures exceeding 1000 Celsius.
    • Calcium bicarbonate is created when it dissolves in water that contains carbon dioxide.

Magnesium Carbonate

It is found in nature as magnesite, which is isomorphic to calcite.

  • Properties: Similar to calcium carbonate in terms of characteristics.

Bicarbonates

  • Sodium Bicarbonates
    • Preparation: By absorption of CO2 in Na2CO3 solution.
    • Uses: used as baking powder.
  • Potassium Bicarbonates
  •       Preparation: Same as NaHCO3
  •       Properties: Same as NaHCO3 but more alkaline and soluble in water.

Chlorides

  • Sodium Chloride
    • Preparation: It is prepared by the method of brine 
    • Uses: To melt ice and snow from the road.
  • Potassium Chloride 
    • Preparation: It occurs in nature as sylvite
    •  Uses: As a fertiliser

Magnesium Chloride

  • Preparation: it is prepared by dissolving MgCO3 in dilute hydrochloric acid. 
  • Properties: It is a deliquescent solid that’s mixed with magnesium oxide to form Sorel cement.

Calcium Chloride

  • Properties: It is a by-product of the Solvay Process and is made of deliquescent crystals.

Sulphates

Sodium Sulphate

  • Preparation: Created by heating common salt with sulphuric acid 
  • Properties: Reduced to Na2S when it is fused with carbon and used in pharmaceutical industries.

Potassium Sulphate

It is found in stassfurt potash deposits such as schonite and kainite and is made by crystallisation.

Magnesium Sulphate

  • Preparation: It is called Epsom salt and obtained by dissolving magnesite in hot dilute H2SO4

Class 11 S Block Elements Notes for Chemistry Chapter 10 

About s-Block Elements Revision Notes

Chapter 10 of Class 11 Chemistry explains the properties, trends, characteristics and importance of the s-Block elements. The Revision Notes for s-Block elements are provided by Extramarks Class 11 Chemistry Chapter 10 Notes which is a summarised version of all the important topics in this chapter. It is written in an easy-to-understand language by the subject matter experts. It is a thoroughly researched material made as per the CBSE examination guidelines. When students study from it, they will get an edge over their peers. To access the resources at Extramarks, students may sign up at Extramarks.

 The topics given below are all covered by Extramarks Notes.

Periodic Trends in the Properties of s-block Elements

  • Alkali Metals
  • Alkali Earth Metals
  • Diagonal Relationship

What are the s-block Elements?

The s-block of the periodic table contains Group 1 and Group 2 elements.

Sub-Topics of the s- Block Elements

  • Anomalous Behaviour of Lithium 
  • Beryllium, Magnesium, and Calcium
  • Characteristics of the Compounds of Alkali Earth Metals 
  • Characteristics of the Compounds of Alkali Metals
  • Group 1 Elements, Alkali Metals
  • Group 2 Elements, Alkali Earth Metals
  • Some Important Compounds of Potassium and Sodium 
FAQs (Frequently Asked Questions)

1. What are some examples of s-Block elements?

The s-block consists of 14 elements. These are Hydrogen (H), Lithium (Li), Helium (He), Sodium (Na), Beryllium (Be), Potassium (K), Magnesium (Mg), Rubidium (Rb), Calcium (Ca), Caesium (Cs), Strontium (Sr), Francium (Fr), Barium (Ba) and Radium (Ra).

2. What are the physical properties of alkali metals?

Alkali Metals have low density and it increases as the group progresses. They generally have a low MP and are soft, light and white and can be easily cut with a knife.