Class 11 Notes

Class 11 Chemistry Notes Chapter 3 (The s- Block Elements) – Chemistry Part-II Book

Philoid

09 May 2025 — 12 min read

Alright class, let's get straight into the s-Block Elements from your NCERT Class 11 Chemistry Part-II book. This chapter is crucial not just for your Class 11 exams but forms a strong foundation for various government exams where chemistry is a component. We'll cover the key concepts, trends, and important compounds in detail.

Chapter 3: The s-Block Elements

Introduction:

The s-block elements are those in which the last electron enters the outermost s-orbital.
They comprise Group 1 (Alkali Metals) and Group 2 (Alkaline Earth Metals).
Their general outer electronic configuration is ns¹ for Group 1 and ns² for Group 2, where 'n' is the outermost principal quantum number.
They are highly reactive metals due to their low ionization enthalpies and large size.
They generally form ionic compounds, except for Lithium (Li) and Beryllium (Be) which show some covalent character.

Group 1: The Alkali Metals

Elements: Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Caesium (Cs), Francium (Fr) (Radioactive).
Electronic Configuration: [Noble Gas] ns¹
Occurrence: Highly reactive, hence not found free in nature. Found as halides, silicates, nitrates, borates, etc. (e.g., NaCl, KCl, NaNO₃). Francium is radioactive with a very short half-life.

General Characteristics & Trends:

Atomic and Ionic Radii: Largest in their respective periods. Increase down the group (Li < Na < K < Rb < Cs) due to the addition of a new shell. Cations (M⁺) are smaller than parent atoms.
Ionization Enthalpy (IE): Have the lowest IE in their respective periods due to large size and only one electron in the valence shell. IE decreases down the group (Li > Na > K > Rb > Cs) as atomic size increases and shielding effect becomes more pronounced. Second IE is very high as the electron needs to be removed from a stable noble gas configuration.
Hydration Enthalpy: Decreases down the group as ionic size increases (Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺). Li⁺ has the maximum degree of hydration due to its smallest size and highest charge density. This is why Li salts are often hydrated (e.g., LiCl·2H₂O) while other alkali metal chlorides are not.
Physical Properties:
- Silvery-white, soft metals (can be cut with a knife). Softness increases down the group.
- Low density, increasing down the group (K is lighter than Na - exception due to unusual increase in atomic size). Li, Na, K are lighter than water.
- Low melting and boiling points due to weak metallic bonding (only one valence electron). Decrease down the group.
- Flame Coloration: Impart characteristic colours to the flame due to easy excitation of the outer electron to higher energy levels. When the electron returns to the ground state, it emits radiation in the visible region.
  - Li: Crimson Red
  - Na: Golden Yellow
  - K: Pale Violet (Lilac)
  - Rb: Red Violet
  - Cs: Blue Violet
- Highly electropositive character.
- Show photoelectric effect (especially K, Rb, Cs) due to low IE. Cs is used in photoelectric cells.
Chemical Properties:
- Highly reactive due to low IE. Reactivity increases down the group.
- Oxidation State: Exhibit +1 oxidation state exclusively in their compounds.
- Reducing Nature: Strong reducing agents (easily lose ns¹ electron). Standard electrode potential (E°) values are highly negative. Li is the strongest reducing agent in aqueous solution despite having the highest IE, due to its very high hydration enthalpy. Reducing power generally increases down the group in gaseous state, but in solution, Li > Cs > Rb > K > Na.
- Reaction with Air/Oxygen: Tarnish rapidly in dry air. Burn vigorously in oxygen.
  - Li forms monoxide (Li₂O) predominantly.
  - Na forms peroxide (Na₂O₂) predominantly (also some Na₂O).
  - K, Rb, Cs form superoxides (KO₂, RbO₂, CsO₂) predominantly.
  - Oxides are basic, basicity increases down the group.
- Reaction with Water: React violently (often explosively) with water to form hydroxides (strong bases) and liberate hydrogen gas (H₂). Reactivity increases down the group.
  - 2M + 2H₂O → 2MOH + H₂ (M = Alkali metal)
- Reaction with Hydrogen: React with H₂ at high temperatures (~673K, Li at 1073K) to form ionic hydrides (M⁺H⁻). Ionic character increases down the group. Hydrides are strong reducing agents.
  - 2M + H₂ → 2MH
- Reaction with Halogens: React vigorously to form ionic halides (M⁺X⁻). Reactivity decreases down the group for a given halogen. LiX shows some covalent character.
  - 2M + X₂ → 2MX
- Solutions in Liquid Ammonia: Dissolve in liquid ammonia to give deep blue, conducting, and paramagnetic solutions.
  - M + (x+y)NH₃ → [M(NH₃)ₓ]⁺ + [e(NH₃)y]⁻ (Ammoniated cation and ammoniated electron)
  - The blue colour is due to the ammoniated electron which absorbs energy in the visible region.
  - The solution is paramagnetic due to the unpaired ammoniated electron.
  - On standing, the blue solution slowly liberates H₂ forming amides: 2M + 2NH₃ → 2MNH₂ + H₂
  - At high concentrations (>3M), the solution becomes bronze-coloured and diamagnetic due to the formation of electron pairs.

Anomalous Behavior of Lithium:

Due to: (i) exceptionally small size of atom and ion, (ii) high polarizing power (charge/radius ratio), (iii) high IE and electronegativity compared to others, (iv) absence of d-orbitals.
Differences from other Alkali Metals:
- Harder, higher melting/boiling point.
- Least reactive but strongest reducing agent in solution.
- Forms monoxide (Li₂O) with O₂, not peroxide or superoxide.
- Forms nitride (Li₃N) on reaction with N₂ (only one in the group).
- LiCl is deliquescent and crystallizes as hydrate (LiCl·2H₂O). Other alkali chlorides are anhydrous.
- LiHCO₃ is not obtained in solid form, while others form solid bicarbonates.
- LiNO₃ decomposes to give Li₂O, NO₂, O₂. Other alkali nitrates decompose to give nitrites (MNO₂) and O₂.
- LiF and Li₂O are comparatively much less soluble in water than corresponding compounds of other alkali metals.
- Shows covalent character in compounds (LiCl soluble in ethanol).
Diagonal Relationship with Magnesium (Mg): Li shows similarities to Mg (Group 2).
- Both are harder and lighter than other metals in their respective groups.
- React slowly with water. Oxides/hydroxides are less soluble.
- Form nitrides (Li₃N, Mg₃N₂).
- Oxides (Li₂O, MgO) do not form peroxides/superoxides easily.
- Carbonates decompose on heating to form oxides and CO₂.
- Solid bicarbonates are not formed.
- Chlorides are deliquescent, hydrated, and soluble in ethanol (covalent nature).

Important Compounds of Sodium:

Sodium Carbonate (Washing Soda, Na₂CO₃·10H₂O):
- Preparation: Solvay Process (Ammonia-Soda Process). Based on the low solubility of NaHCO₃ in ammoniacal brine.
  - 2NH₃(aq) + H₂O(l) + CO₂(g) → (NH₄)₂CO₃(aq)
  - (NH₄)₂CO₃(aq) + H₂O(l) + CO₂(g) → 2NH₄HCO₃(aq)
  - NH₄HCO₃(aq) + NaCl(aq) → NaHCO₃(s) + NH₄Cl(aq) (NaHCO₃ precipitates)
  - 2NaHCO₃(s) --Heat--> Na₂CO₃(s) + CO₂(g) + H₂O(g)
  - Ammonia recovery: 2NH₄Cl(aq) + Ca(OH)₂(aq) --Heat--> 2NH₃(g) + CaCl₂(aq) + 2H₂O(l)
  - Note: K₂CO₃ cannot be prepared by Solvay process because KHCO₃ is too soluble.
- Properties: White crystalline solid, efflorescent (loses water of crystallization), soluble in water (basic solution due to hydrolysis), melts on heating.
- Uses: Water softening, laundering, cleaning, manufacture of glass, soap, borax, caustic soda, paper, paints.
Sodium Chloride (Common Salt, NaCl):
- Occurrence: Seawater (2.7-2.9% by mass), salt lakes, rock salt deposits.
- Properties: White crystalline solid, high melting point, soluble in water, slightly soluble in ethanol. Crude salt is hygroscopic due to impurities like MgCl₂ and CaCl₂.
- Uses: Essential dietary component, preservative, source for manufacturing NaOH, Na₂CO₃, Cl₂, H₂, Na.
Sodium Hydroxide (Caustic Soda, NaOH):
- Preparation: Castner-Kellner Process (Electrolysis of brine - NaCl solution using a mercury cathode and carbon anode).
  - Cathode (Hg): Na⁺ + e⁻ → Na-Hg (Sodium amalgam)
  - Anode (C): Cl⁻ → ½ Cl₂ + e⁻
  - Amalgam treated with water: 2Na-Hg + 2H₂O → 2NaOH + 2Hg + H₂
- Properties: White, translucent, deliquescent solid. Highly soluble in water (exothermic), strong base. Absorbs CO₂ from air (forms Na₂CO₃). Reacts with acids, non-metals, metals (like Al, Zn).
- Uses: Manufacture of soap, paper, artificial silk (rayon), dyes, petroleum refining, purification of bauxite, laboratory reagent.
Sodium Bicarbonate (Baking Soda, NaHCO₃):
- Preparation: Saturating a solution of Na₂CO₃ with CO₂. Or obtained as intermediate in Solvay process.
  - Na₂CO₃ + H₂O + CO₂ → 2NaHCO₃
- Properties: White crystalline powder, sparingly soluble in water. Decomposes on heating (~373K) to give Na₂CO₃, CO₂, H₂O. Mildly basic.
- Uses: Baking (releases CO₂ making cakes fluffy), mild antiseptic, fire extinguishers, laboratory reagent.

Biological Importance of Na and K:

Found primarily in body fluids (Na⁺ outside cells, K⁺ inside cells).
Maintain osmotic pressure, water balance.
Involved in nerve impulse transmission (Sodium-Potassium Pump).
K⁺ activates many enzymes.

Group 2: The Alkaline Earth Metals

Elements: Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba), Radium (Ra) (Radioactive).
Electronic Configuration: [Noble Gas] ns²
Occurrence: Reactive, found combined in nature. E.g., Mg (Carnallite, Dolomite, Epsomite), Ca (Limestone, Gypsum, Dolomite), Sr (Celestite), Ba (Barytes). Radium is rare, found in uranium ores.

General Characteristics & Trends:

Atomic and Ionic Radii: Smaller than corresponding alkali metals due to increased nuclear charge. Increase down the group (Be < Mg < Ca < Sr < Ba). Cations (M²⁺) are smaller than parent atoms.
Ionization Enthalpy (IE): Higher than alkali metals due to smaller size and increased nuclear charge. IE decreases down the group (Be > Mg > Ca > Sr > Ba). Second IE (IE₂) is higher than first IE (IE₁), but the sum (IE₁ + IE₂) is low enough for them to form M²⁺ ions.
Hydration Enthalpy: Higher than alkali metals due to smaller size and higher charge (+2). Decreases down the group as ionic size increases (Be²⁺ > Mg²⁺ > Ca²⁺ > Sr²⁺ > Ba²⁺). Be²⁺ has very high hydration enthalpy. Hydrated salts are common (e.g., MgCl₂·6H₂O, CaCl₂·6H₂O).
Physical Properties:
- Silvery-white, lustrous, harder than alkali metals (stronger metallic bonding due to two valence electrons). Hardness decreases down the group.
- Higher melting and boiling points than alkali metals (no clear trend). Be and Mg are greyish.
- Higher densities than alkali metals. Density trend is irregular (Ca < Mg < Be < Sr < Ba).
- Flame Coloration: Impart characteristic colours (except Be and Mg - due to high IE, electrons are too strongly bound).
  - Ca: Brick Red
  - Sr: Crimson Red
  - Ba: Apple Green
  - Ra: Crimson
- Electropositive, but less than alkali metals.
Chemical Properties:
- Less reactive than alkali metals. Reactivity increases down the group (Be < Mg < Ca < Sr < Ba).
- Oxidation State: Exhibit +2 oxidation state exclusively.
- Reducing Nature: Strong reducing agents, but weaker than alkali metals. Electrode potentials are less negative than alkali metals. Reducing power increases down the group.
- Reaction with Air/Oxygen: Less reactive towards air. Burn in oxygen when heated to form oxides (MO). Also react with nitrogen upon heating to form nitrides (M₃N₂).
  - 2M + O₂ --Heat--> 2MO
  - 3M + N₂ --Heat--> M₃N₂
- Reaction with Water: React with water (less vigorously than alkali metals) to form hydroxides and H₂. Reactivity increases down the group.
  - Be does not react with water.
  - Mg reacts with hot water/steam.
  - Ca, Sr, Ba react readily with cold water.
  - M + 2H₂O → M(OH)₂ + H₂
  - Hydroxides are basic (less basic than alkali metal hydroxides). Basicity increases down the group (Be(OH)₂ amphoteric, Mg(OH)₂ weakly basic, others reasonably strong bases).
- Reaction with Hydrogen: Form hydrides (MH₂) upon heating (except Be). BeH₂ and MgH₂ are covalent and polymeric, others are ionic.
  - M + H₂ --Heat--> MH₂
- Reaction with Halogens: React readily with halogens upon heating to form dihalides (MX₂).
  - M + X₂ --Heat--> MX₂
  - BeF₂ is ionic, but BeCl₂ is covalent (soluble in organic solvents, polymeric structure). Other halides are ionic. Fluorides are less soluble than chlorides.
- Solutions in Liquid Ammonia: Like alkali metals, dissolve in liquid ammonia to give deep blue-black solutions containing ammoniated electrons.
- Solubility and Thermal Stability:
  - Carbonates (MCO₃): Sparingly soluble in water. Solubility decreases down the group. Thermal stability increases down the group (BeCO₃ unstable, BaCO₃ most stable). Decompose on heating: MCO₃ --Heat--> MO + CO₂.
  - Sulphates (MSO₄): BeSO₄ and MgSO₄ are readily soluble (high hydration enthalpy). Solubility decreases down the group (CaSO₄ sparingly soluble, SrSO₄, BaSO₄ almost insoluble). Thermal stability increases down the group.
  - Hydroxides (M(OH)₂): Solubility increases down the group (opposite to carbonates/sulphates). Be(OH)₂ and Mg(OH)₂ almost insoluble, Ba(OH)₂ fairly soluble. Thermal stability increases down the group. Basicity increases down the group.
  - Nitrates (M(NO₃)₂): Decompose on heating: 2M(NO₃)₂ --Heat--> 2MO + 4NO₂ + O₂ (like LiNO₃).

Anomalous Behavior of Beryllium:

Due to: (i) very small size, (ii) high IE and electronegativity, (iii) high polarizing power, (iv) absence of d-orbitals.
Differences from other Alkaline Earth Metals:
- Hardest metal in the group.
- Highest melting/boiling point.
- Forms covalent compounds predominantly (due to high IE and small size).
- Does not react with water even at high temperatures.
- Oxide (BeO) and Hydroxide (Be(OH)₂) are amphoteric (react with both acids and bases). Others are basic.
- Does not exhibit coordination number more than 4 (no d-orbitals). Others can have C.N. = 6.
- Carbide (Be₂C) gives methane on hydrolysis, others give acetylene or other hydrocarbons.
- Does not react directly with hydrogen. BeH₂ prepared indirectly.
Diagonal Relationship with Aluminium (Al): Be shows similarities to Al (Group 13).
- Both metals are resistant to acid attack due to a protective oxide film.
- Both react with strong alkalis liberating H₂ (Be(OH)₂ and Al(OH)₃ are amphoteric).
  - Be + 2NaOH + 2H₂O → Na₂[Be(OH)₄] + H₂
  - 2Al + 2NaOH + 6H₂O → 2Na[Al(OH)₄] + 3H₂
- Chlorides (BeCl₂, AlCl₃) are covalent, soluble in organic solvents, act as Lewis acids, and are used as Friedel-Crafts catalysts. Both have bridged structures in vapour phase/solid state.
- Both form fluoro complexes ([BeF₄]²⁻, [AlF₆]³⁻).
- Oxides (BeO, Al₂O₃) are hard, high melting solids.

Important Compounds of Calcium:

Calcium Oxide (Quick Lime, CaO):
- Preparation: Heating limestone (CaCO₃) strongly (~1070-1270 K) in a rotary kiln. Reaction is reversible, CO₂ must be removed.
  - CaCO₃(s) ⇌ CaO(s) + CO₂(g)
- Properties: White amorphous solid, high melting point (~2870 K). Basic oxide. Reacts exothermically with water ('slaking of lime') to form Ca(OH)₂. Reacts with acidic oxides (like SiO₂) and acids. Absorbs moisture and CO₂ from air.
  - CaO + H₂O → Ca(OH)₂ + Heat
  - CaO + SiO₂ → CaSiO₃ (Calcium silicate - slag)
  - CaO + CO₂ → CaCO₃
- Uses: Manufacture of cement, mortar, glass, calcium carbide, sodium carbonate (from caustic soda), purification of sugar, drying agent (basic), steel making (flux).
Calcium Hydroxide (Slaked Lime, Ca(OH)₂):
- Preparation: Adding water to quick lime (CaO).
- Properties: White amorphous powder, sparingly soluble in water. Aqueous solution is 'lime water', suspension is 'milk of lime'. Reacts with CO₂ to form CaCO₃ (lime water turns milky), excess CO₂ forms soluble calcium bicarbonate Ca(HCO₃)₂ (milkiness disappears). Reacts with Cl₂ to form bleaching powder Ca(OCl)₂·CaCl₂·Ca(OH)₂·2H₂O. Basic.
  - Ca(OH)₂ + CO₂ → CaCO₃(s) + H₂O
  - CaCO₃(s) + H₂O + CO₂(excess) → Ca(HCO₃)₂(aq)
  - 2Ca(OH)₂ + 2Cl₂ → CaCl₂ + Ca(OCl)₂ + 2H₂O (Bleaching powder components)
- Uses: Mortar (building material), whitewashing, glass making, tanning industry, sugar purification, bleaching powder manufacture, softening of hard water.
Calcium Carbonate (Limestone, Marble, Chalk, CaCO₃):
- Occurrence: Abundant in nature in various forms.
- Preparation: Passing CO₂ through slaked lime or adding Na₂CO₃ to CaCl₂ solution.
  - Ca(OH)₂ + CO₂ → CaCO₃ + H₂O
  - CaCl₂ + Na₂CO₃ → CaCO₃ + 2NaCl
- Properties: White fluffy powder or crystalline solid, almost insoluble in water. Decomposes on heating above ~1200 K. Reacts with dilute acids to liberate CO₂.
  - CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂
- Uses: Building material (marble), manufacture of quick lime, cement, flux in metallurgy (removes impurities), manufacture of high-quality paper, antacid, toothpaste abrasive, chewing gum filler.
Calcium Sulphate (Plaster of Paris, PoP, CaSO₄·½H₂O):
- Preparation: Heating Gypsum (CaSO₄·2H₂O) carefully to ~393 K (120°C). If heated above 393 K, anhydrous CaSO₄ ('dead burnt plaster') is formed, which sets slowly.
  - 2(CaSO₄·2H₂O) --~393K--> 2(CaSO₄)·H₂O + 3H₂O (or written as CaSO₄·½H₂O)
- Properties: White powder. On mixing with water (1/3rd its mass), it forms a plastic mass that sets into a hard solid (due to re-formation of gypsum) within 5-15 minutes, with slight expansion.
  - CaSO₄·½H₂O + 1½H₂O → CaSO₄·2H₂O (Setting of Plaster)
- Uses: Making casts for statues, busts, decorative work, surgical bandages for setting fractured bones, dentistry, building material (false ceilings).
Cement: Complex mixture. Portland cement is most common.
- Composition: CaO (50-60%), SiO₂ (20-25%), Al₂O₃ (5-10%), MgO (2-3%), Fe₂O₃ (1-2%), SO₃ (1-2%). Raw materials: Limestone (CaCO₃) and Clay (containing SiO₂, Al₂O₃, Fe₂O₃). Good cement ratio: SiO₂/Al₂O₃ = 2.5-4; CaO/(SiO₂ + Al₂O₃ + Fe₂O₃) ≈ 2.
- Manufacture: Strongly heating limestone and clay mixture in a rotary kiln. 'Cement clinker' formed. Clinker mixed with 2-3% gypsum (CaSO₄·2H₂O) and ground to fine powder. Gypsum slows down the setting process.
- Setting: When mixed with water, hydration of constituents occurs, forming a hard mass (interlocking crystals of hydrates). Complex reactions involving hydration and hydrolysis of calcium silicates and aluminates.

Biological Importance of Mg and Ca:

Mg²⁺: Concentrated in animal cells, cofactor for enzymes using ATP (phosphate transfer), main pigment in chlorophyll (photosynthesis).
Ca²⁺: Major component of bones and teeth (as apatite), important in blood clotting, muscle contraction, nerve function, cell membrane integrity.

This covers the essential aspects of the s-Block elements relevant for your exams. Focus on the trends, reasons for anomalous behavior, diagonal relationships, properties and preparation methods (especially named processes like Solvay, Castner-Kellner) of important compounds, and their uses.

Multiple Choice Questions (MCQs):

Which alkali metal is the strongest reducing agent in aqueous solution?
(a) Na
(b) K
(c) Li
(d) Cs
The correct order of decreasing second ionization enthalpy for Group 1 elements is:
(a) Li > Na > K > Rb > Cs
(b) Cs > Rb > K > Na > Li
(c) Li > Cs > Rb > K > Na
(d) Order cannot be predicted easily
Which of the following alkaline earth metal hydroxides is amphoteric in nature?
(a) Ca(OH)₂
(b) Sr(OH)₂
(c) Be(OH)₂
(d) Mg(OH)₂
When sodium metal is dissolved in liquid ammonia, the deep blue colour is attributed to:
(a) Ammoniated Na⁺ ions
(b) Ammoniated electrons
(c) Formation of sodium amide (NaNH₂)
(d) Formation of sodium hydride (NaH)
The Solvay process is used for the manufacture of:
(a) NaOH
(b) Na₂CO₃
(c) NaCl
(d) NaHCO₃ (as final product)
Which property decreases down the group for Alkaline Earth Metals?
(a) Atomic Radii
(b) Basicity of Hydroxides
(c) Solubility of Sulphates in water
(d) Reactivity with water
Beryllium shows a diagonal relationship with:
(a) Magnesium (Mg)
(b) Aluminium (Al)
(c) Boron (B)
(d) Silicon (Si)
The formula for Plaster of Paris is:
(a) CaSO₄·2H₂O
(b) CaSO₄·H₂O
(c) 2CaSO₄·H₂O (or CaSO₄·½H₂O)
(d) CaSO₄
Which element imparts a brick-red colour to the Bunsen flame?
(a) Na
(b) K
(c) Ca
(d) Ba
Which of the following is NOT a use of Calcium Carbonate (CaCO₃)?
(a) Manufacture of Cement
(b) Antacid
(c) Setting fractured bones
(d) Flux in metallurgy

Answers:

(c) Li (due to high hydration enthalpy)
(a) Li > Na > K > Rb > Cs (Follows the trend of first IE as electron is removed from M⁺ ion)
(c) Be(OH)₂
(b) Ammoniated electrons
(b) Na₂CO₃
(c) Solubility of Sulphates in water (BeSO₄, MgSO₄ soluble; BaSO₄ insoluble)
(b) Aluminium (Al)
(c) 2CaSO₄·H₂O (or CaSO₄·½H₂O)
(c) Ca
(c) Setting fractured bones (This is a use of Plaster of Paris)

Study these notes thoroughly, focusing on understanding the trends and exceptions. Good luck with your preparation!

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