s-Block

s-Block

🟢 Lite — Quick Review (1h–1d)

Rapid summary for last-minute revision before your exam.

The s-block comprises the two left-most groups of the periodic table whose differentiating electron enters the outermost s-orbital: Group 1 alkali metals (ns¹) — Li, Na, K, Rb, Cs, Fr — and Group 2 alkaline earth metals (ns²) — Be, Mg, Ca, Sr, Ba, Ra.

Reactivity with water defines the families: 2M + 2H₂O → 2MOH + H₂ (alkali metals, increasingly violent down the group) and M + 2H₂O → M(OH)₂ + H₂ (alkaline earths, far less vigorous). Ionisation enthalpy decreases down both groups, making heavier members the strongest reducing agents in their respective columns.

Three high-yield pointers for CUET UG:

• Li and Be are anomalous — small size, high charge density, and a diagonal relationship with Mg and Al respectively. BeO and Be(OH)₂ are amphoteric, not basic.
• Oxidation product of the metal in O₂ depends on size: Li → Li₂O (normal oxide), Na → Na₂O₂ (peroxide), K/Rb/Cs → MO₂ (superoxide).
• Flame colours are diagnostic: Na = yellow, K = violet (viewed through blue cobalt glass), Ca = brick red, Sr = crimson, Ba = apple green.

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🟡 Standard — Regular Study (2d–2mo)

Standard content for students with a few days to months.

Electronic Configuration and Position

Every s-block atom ends with ns¹ (Group 1) or ns² (Group 2). Because the outermost shell holds only one or two electrons and these sit far from the nucleus in heavy elements, the group-1 and group-2 cations achieve noble-gas configurations by losing just one or two electrons, which is why both families are fiercely electropositive.

Periodic Trends

Across a period (e.g. Na → Mg), atomic and ionic radii shrink because the same nuclear charge pulls a smaller shell inward. Down a group, radii grow sharply as each successive shell is added. Consequences:

• Ionisation enthalpy (IE) drops down each group — Cs has the lowest first IE of any stable element.
• Metallic character, electropositivity and reducing power all increase down each group, so Cs is the strongest metallic reducing agent, Li the weakest among alkali metals.
• Hydration enthalpy (ΔH_hyd) follows ΔH_hyd ∝ (z⁺·z⁻)/(r⁺ + r⁻); small Li⁺ is heavily hydrated, large Cs⁺ weakly, producing the famous abnormal molar-conductivity order Λ°m(Li⁺) < Λ°m(Na⁺) < Λ°m(K⁺) < Λ°m(Rb⁺) < Λ°m(Cs⁺) — the size trend wins.

Anomalous Behaviour and Diagonal Relationships

Property	Li / Be anomaly	Diagonal partner
Oxide / hydroxide	Be(OH)₂ amphoteric	Al(OH)₃
Carbonate	Li₂CO₃ decomposes easily	MgCO₃
Reaction with O₂	Li → Li₂O only	Mg → MgO only

The high charge/radius ratio polarises anion electron clouds, producing covalent character in Li–X and Be–X bonds.

Compounds of Industrial Importance

Plaster of Paris is CaSO₄·½H₂O, prepared by heating gypsum to 393 K; it sets to gypsum on mixing with water. Bleaching powder is CaOCl₂, made by passing Cl₂ over dry slaked lime — it is the standard chlorinating agent for water treatment. Na₂O₂ is used in submarines and self-contained breathing apparatus because it liberates O₂ on reacting with exhaled CO₂.

CUET Question Patterns

Expect 2–3 MCQs each year drawn from: flame test identification of Group 1 and Group 2 cations; matching the metal oxide / peroxide / superoxide; reasoning about solubility trends using lattice energy vs hydration energy balance (Li₂CO₃ and BeCO₃ are insoluble despite being Group 1/2 carbonates, because their lattice enthalpy dominates).

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🔴 Extended — Deep Study (3mo+)

Comprehensive coverage for students on a longer study timeline.

Solubility Decisions: Lattice vs Hydration Enthalpy

The solubility of an ionic salt ≈ balance of lattice enthalpy (ΔH_lat) and hydration enthalpy (ΔH_hyd). Both magnitudes decrease down the group, but at different rates:

• For hydroxides, ΔH_hyd falls faster than ΔH_lat, so solubility increases down Group 1 (NaOH is more soluble than LiOH). Among Group 2, Ca(OH)₂ is moderately soluble (used in whitewash), Sr(OH)₂ and Ba(OH)₂ are clearly soluble.
• For carbonates and sulphates, ΔH_lat dominates, hence solubility decreases down the group. Li₂CO₃ and BeCO₃ are sparingly soluble — the same trap that catches students in NEET also features in CUET.

Thermal Stability of Carbonates and Nitrates

MCO₃ → MO + CO₂ becomes harder to decompose down Group 2 because the larger M²⁺ polarises the CO₃²⁻ less effectively; thus BeCO₃ decomposes at ~373 K, while BaCO₃ requires strong heating. Nitrates behave analogously: 2M(NO₃)₂ → 2MO + 4NO₂ + O₂, with the temperature rising markedly from Be to Ba. Group 1 nitrates give MNO₂ + ½O₂ on heating; ammonium nitrate routes are not relevant here.

Edge Cases and Common Traps

• BeCl₂ is polymeric and covalent in the solid (bridge-Cl structure), turning monomeric and linear (sp) only in the gas phase above 1200 K. Statements claiming it is purely ionic are wrong.
• Na₂O₂ vs Na₂O vs KO₂: peroxide anion is O₂²⁻ (oxidation state of O = −1); superoxide is O₂⁻ (O = −½). Memorise the colour — peroxides are pale yellow, superoxides are yellow-to-orange.
• Cs and Fr are radioactive in any meaningful quantity; CUET rarely asks beyond mentioning their position.
• Magnesium is passivated by a thin MgO layer in cold water and reacts vigorously only on heating — students forget this and rank Mg “more reactive than Na”, which is false (Na reacts violently in cold water).

Worked Numeric Hook

Compare dissolution of LiF and CsF. Although both are alkali halides, LiF has hydration enthalpy ≈ −1034 kJ mol⁻¹ and lattice enthalpy ≈ −1036 kJ mol⁻¹, giving a near-zero net ΔH — LiF is poorly soluble. CsF has ΔH_lat ≈ −636 kJ mol⁻¹ but ΔH_hyd ≈ −667 kJ mol⁻¹, so dissolution is exothermic and CsF is highly soluble. This single calculation explains why fluorides invert the usual “all Group 1 salts are soluble” claim.

Practice Prompts

1. Identify the products when Rb is burnt in excess O₂ and write the balanced equation. Why is this product used in closed-circuit breathing apparatus?
2. Arrange Be(OH)₂, Mg(OH)₂, Ca(OH)₂, Sr(OH)₂, Ba(OH)₂ in order of increasing solubility in water and justify using enthalpy data.

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