s-Block

🟢 Lite — Quick Review (1h–1d)

Rapid summary for last-minute revision before your exam.

The s-block elements are those in which the outermost electron enters an s-orbital. They consist of Group 1 (alkali metals: Li, Na, K, Rb, Cs, Fr) and Group 2 (alkaline earth metals: Be, Mg, Ca, Sr, Ba, Ra). They are characterised by having ns¹ and ns² outer electronic configurations respectively. The s-block lies on the far left of the periodic table.

Electronic Configurations:

Element	Symbol	Atomic No.	Configuration
Lithium	Li	3	[He] 2s¹
Sodium	Na	11	[Ne] 3s¹
Potassium	K	19	[Ar] 4s¹
Rubidium	Rb	37	[Kr] 5s¹
Caesium	Cs	55	[Xe] 6s¹
Francium	Fr	87	[Rn] 7s¹ (radioactive)
Beryllium	Be	4	[He] 2s²
Magnesium	Mg	12	[Ne] 3s²
Calcium	Ca	20	[Ar] 4s²
Strontium	Sr	38	[Kr] 5s²
Barium	Ba	56	[Xe] 6s²
Radium	Ra	88	[Rn] 7s² (radioactive)

Key Properties:

Property	Group 1 (Alkali)	Group 2 (Alkaline Earth)
Valence electrons	ns¹	ns²
Oxidation state	+1	+2
Ionisation energy	Low (lowest in period)	Higher than Group 1
Atomic radius	Larger than Group 2 in same period	Smaller than Group 1
Hydration enthalpy	Very high (small ions)	Very high (2+ ions)
Flame test	Golden yellow (Na), lilac (K)	Brick red (Ca), apple green (Ba)
Nature of oxides	Basic/strongly basic	Basic (less than Group 1)

⚡ Exam Tips:

• Na and K are essential for nerve impulse transmission (Na⁺/K⁺ ATPase pump)
• Mg is central to chlorophyll (porphyrin ring with Mg²⁺); Ca is central to bones and teeth (hydroxyapatite: Ca₁₀(PO₄)₆(OH)₂)
• BeCl₂ is covalent; other Group 2 chlorides are ionic
• Solubility of hydroxides increases down the group: Be(OH)₂ < Mg(OH)₂ < Ca(OH)₂ < Sr(OH)₂ < Ba(OH)₂
• Thermal stability of carbonates: BeCO₃ < MgCO₃ < CaCO₃ < SrCO₃ < BaCO₃ (decreases down group = easier to decompose)

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🟡 Standard — Regular Study (2d–2mo)

Standard content for students with a few days to months.

Physical Properties of s-Block Elements:

Atomic and Ionic Radii: Atomic radii increase down the group (more electron shells). For isoelectronic species (same number of electrons), the cation is smaller than the neutral atom, and the anion is larger.

• Na⁺ (102 pm) < Mg²⁺ (72 pm) — both have Ne configuration; Mg²⁺ is smaller due to higher nuclear charge
• The anomalous properties of Li and Be are due to their small size — Li and Be are atypical; their properties resemble diagonal neighbours (Mg and Al respectively)

Ionisation Enthalpy: Ionisation enthalpy decreases down the group as atomic size increases and shielding increases.

• First IE₂ << IE₁ for Group 1 (removing one electron gives stable octet)
• For Group 2, IE₁ is higher than Group 1 because of higher nuclear charge and smaller radius
• IE₂ is much higher than IE₁ for Group 2 because the second electron must be removed from a stable octet configuration

Hydration Enthalpy: M²⁺ ions (Group 2) have higher charge density than M⁺ ions (Group 1), so they have more negative (stronger) hydration enthalpies. However, this also means Group 2 ions form stronger ion pairs and their compounds are more lattice-dominated.

Flame Colours: When metal ions are heated, electrons are excited to higher energy levels. As they return, they emit photons of characteristic wavelengths:

Ion	Flame Colour
Li⁺	Crimson red
Na⁺	Golden yellow (persistent — detect traces at 0.00001%)
K⁺	Lilac (pale violet —透过钴玻璃观看)
Ca²⁺	Brick red
Sr²⁺	Crimson red
Ba²⁺	Apple green

The flame test is a classic qualitative analysis technique for detecting Group 1 and 2 metals.

Anomalous Behaviour of Lithium and Beryllium:

Lithium (Diagonal relationship with Magnesium):

• Li is the only Group 1 element that forms a nitride: 6Li + N₂ → 2Li₃N
• Li₂O is formed directly (not peroxide like Na): 4Li + O₂ → 2Li₂O
• LiOH decomposes at high T: 2LiOH → Li₂O + H₂O (unlike other alkali hydroxides which are stable)
• Li forms complex ions less readily; is the most covalent of alkali metals
• Solubility: LiF and Li₂CO₃ are sparingly soluble (unlike other alkali metal salts which are soluble)

Beryllium (Diagonal relationship with Aluminium):

• BeCl₂ is covalent (forms chloroberyllate ion [BeCl₄]²⁻); MgCl₂ is ionic
• Be forms complexes (e.g., [BeF₃]⁻) — other Group 2 do not
• Be does not react with water (unlike Mg, Ca, Sr, Ba); Al also doesn’t react with cold water
• BeO is amphoteric (dissolves in acids AND alkalies); forms [Be(OH)₄]²⁻ with NaOH; MgO is basic only
• Be has high ionisation energy and small size — doesn’t form simple cations easily

Important Compounds:

Sodium Hydroxide (NaOH — Caustic Soda):

• Manufactured by Castner-Kellner process (mercury cell)
  • Mercury cathode: Na⁺ + e⁻ → Na (dissolves in mercury → Na amalgam)
  • Mercury anode: 2Cl⁻ → Cl₂
  • Na amalgam reacts with water: 2Na + 2H₂O → 2NaOH + H₂
• Strong base; absorbs CO₂ from air: 2NaOH + CO₂ → Na₂CO₃ + H₂O
• White, deliquescent solid; dissolves in water with much heat

Sodium Carbonate (Na₂CO₃ — Washing Soda):

• Manufactured by Solvay process (NH₃ + CO₂ + NaCl + H₂O)
• Steps:
  1. NH₃ + CO₂ + H₂O → NH₄HCO₃
  2. NH₄HCO₃ + NaCl → NaHCO₃↓ + NH₄Cl
  3. 2NaHCO₃ → Na₂CO₃ + CO₂ + H₂O (heated —“soda ash”)
• The CO₂ from step 3 is recycled
• By-product NH₄Cl is a fertilizer or recovered
• Also known as soda ash; used in glass, soap, and water softening

Quick Lime (CaO) and Slaked Lime (Ca(OH)₂):

• CaCO₃ → CaO + CO₂ (at 900°C in lime kiln)
• CaO + H₂O → Ca(OH)₂ (vigorous exothermic reaction)
• Used in: steel making (flux), whitewashing (Ca(OH)₂ + CO₂ → CaCO₃), sugar refining, pH adjustment

Plaster of Paris (CaSO₄·0.5H₂O) and Dead Burnt Plaster:

• CaSO₄·2H₂O (gypsum) heated to 373–393 K → CaSO₄·0.5H₂O + 1.5H₂O
• When mixed with water: 2CaSO₄·0.5H₂O + 3H₂O → 2CaSO₄·2H₂O
• Sets hard by crystallisation — used for making casts, moulds, and in building
• Dead burnt plaster: Gypsum heated above 473 K → anhydrous CaSO₄ (does not set with water)

Cement: Cement is manufactured from limestone + clay, heated in rotary kiln at ~1700 K:

• Forms clinker (alite Ca₃SiO₅, belite Ca₂SiO₄, aluminate Ca₃Al₂O₆, ferrite Ca₄Al₂Fe₂O₁₀)
• Ground with 3–5% gypsum (regulates setting time)
• Setting of cement: Complex hydration reactions form interlocking calcium silicate hydrate (C-S-H) crystals

⚡ Common Mistakes:

• Thermal stability of carbonates: BeCO₃ decomposes at low T (100°C); BaCO₃ requires ~1300°C — this trend is opposite to the trend in hydroxides
• Solubility of hydroxides: increases down Group 2 — Be(OH)₂ is insoluble but Ba(OH)₂ is fairly soluble
• Be doesn’t form simple Be²⁺ ions in solution — exists as [Be(H₂O)₄]²⁺ which is highly acidic (pKa ~5.4) and hydrolyses
• Na₂CO₃·10H₂O (washing soda) is different from Na₂CO₃ (soda ash/anhydrous)

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🔴 Extended — Deep Dive (exam-level mastery)

For students preparing for top-rank selection.

Why s-Block Elements Are So Reactive:

Alkali metals have one loosely held s-electron. Their first ionisation energies are among the lowest of all elements:

• Li: 520 kJ/mol; Na: 496; K: 419; Cs: 376 kJ/mol

The low IE means they readily form M⁺ ions. The hydration enthalpy of the resulting ion is very large (exothermic) because the small, highly charged cation strongly polarises water molecules. The lattice energy of the resulting compounds is also high (favourable formation of ionic solids). Overall, the formation of alkali metal salts is energetically favourable.

Alkaline earth metals have two s-electrons. Their second IE is higher than the first, but the overall process M → M²⁺ + 2e⁻ is still favourable because:

1. The M²⁺ hydration enthalpy is very large (more than compensates for double IE)
2. The lattice energy of MX₂ salts is high

Energetics of Ionic Compound Formation (Born-Haber Cycle):

For NaCl: ΔH°f = ΔHsub(Na) + IE₁(Na) + ½ΔHdiss(Cl₂) + EA(Cl) + U (lattice energy)

Where:

• ΔHsub = enthalpy of sublimation of Na (metal)
• IE₁ = first ionisation energy of Na
• ½ΔHdiss = half the bond dissociation energy of Cl₂
• EA = electron affinity of Cl
• U = lattice energy (exothermic, negative)

The lattice energy for NaCl: U = –787 kJ/mol (very large, drives the reaction).

For MgCl₂ (two Cl⁻ ions): ΔH°f = ΔHsub(Mg) + IE₁ + IE₂(Mg) + ΔHdiss(Cl₂) + 2×EA(Cl) + U(MgCl₂)

Diagonal Relationship — Why Does It Exist?

In the periodic table, atomic size decreases across a period and increases down a group. The diagonal relationship occurs between elements where the decrease across a period roughly equals the increase down a group:

• Li (period 2) → size decreases significantly across to Be (covalent character increases)
• Mg (period 3) → similar size to Li

Similarly:

• Be (period 2) → small size, high ionisation, covalent character
• Al (period 3) → small ionic radius, amphoteric oxide, covalent chloride

Covalent Character of s-Block Compounds (Fajans’ Rules):

Fajans’ rules explain why some s-block compounds are covalent:

1. Small cation: Be²⁺ (45 pm) and Li⁺ (76 pm) are small → high charge density → strong polarising power
2. High cation charge: M²⁺ polarises more than M⁺
3. Large anion: Anions with large ionic radius (I⁻, S²⁻) are easily polarised

Consequences:

• BeI₂ is covalent (I⁻ is large, easily polarised)
• MgI₂ is more covalent than MgCl₂
• LiI is more covalent than LiF

Comparison of Solubility Trends:

Group 1 carbonates: All soluble except Li₂CO₃ (sparingly soluble due to strong Li⁺–CO₃²⁻ attraction) Group 2 carbonates: Practically insoluble; solubility decreases down the group (BeCO₃ > MgCO₃ > CaCO₃ > SrCO₃ > BaCO₃)

The lattice energy of BeCO₃ is high (small ions) but hydration enthalpy is relatively low. As ionic size increases down the group, lattice energy decreases more rapidly than hydration enthalpy → overall solubility decreases.

Group 1 hydroxides: All soluble; solubility increases down the group Group 2 hydroxides: Sparingly soluble; solubility increases down the group

• Be(OH)₂ is amphoteric and precipitates at pH 7; Ba(OH)₂ is quite soluble

Nitrates of both groups are all soluble.

Biological Importance:

Element	Role
Na⁺	Resting membrane potential (–70 mV); Na⁺/K⁺-ATPase maintains gradient; nerve impulse propagation
K⁺	Intracellular fluid major cation; essential for protein synthesis; heart rhythm
Mg²⁺	Central atom in chlorophyll; cofactor for ATPases; stabilises DNA; deficiency causes muscle spasms
Ca²⁺	Bones and teeth (hydroxyapatite); muscle contraction (troponin C); blood clotting (Factor IV); intracellular signalling
Ca²⁺/Mg²⁺	Combined → essential for normal heart rhythm, nerve function, muscle contraction

Hypercalcaemia (>2.9 mmol/L): Causes confusion, nausea, cardiac arrhythmias Hypocalcaemia (<2.1 mmol/L): Causes tetany (muscle spasms), seizures, cardiac arrhythmias

Hard Water:

• Temporary hardness: Ca(HCO₃)₂, Mg(HCO₃)₂ — removed by boiling Ca(HCO₃)₂ → CaCO₃↓ + CO₂ + H₂O
• Permanent hardness: CaSO₄, MgSO₄, CaCl₂ — removed by ion exchange (zeolite process) Ca²⁺ (in water) + Na₂Z (zeolite) → CaZ + 2Na⁺ Zeolite is regenerated with NaCl solution

NEET High-Yield Pattern:

• Diagonal: Li ↔ Mg; Be ↔ Al
• Flame: Na = yellow; K = lilac; Ca = brick red; Ba = green
• BeCl₂ is covalent; other G2 chlorides are ionic
• Hydroxide solubility: increases down Group 2
• Carbonate/thermal stability: decreases down Group 2 (opposite to hydroxide)
• Solvay process: NH₃ + CO₂ + NaCl + H₂O → NaHCO₃↓ → Na₂CO₃
• Castner process for NaOH (mercury cell)
• Biological: Na/K pump; Mg in chlorophyll; Ca in bones
• Anomalous Li: forms nitride (Li₃N), oxide (Li₂O), decomposes LiOH
• Anomalous Be: amphoteric oxide (BeO), covalent BeCl₂, no reaction with water