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Chemistry 3% exam weight

s-Block

Part of the JEE Main study roadmap. Chemistry topic chem-015 of Chemistry.

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s-Block

🟢 Lite — Quick Review (1h–1d)

Rapid summary for last-minute revision before your exam.

The s-block houses Group 1 (alkali metals) and Group 2 (alkaline earth metals), whose valence electrons occupy an s-orbital: ns¹ for alkali metals (Li, Na, K, Rb, Cs) and ns² for alkaline earths (Be, Mg, Ca, Sr, Ba). These elements are soft, low-melting metals that readily lose electrons to form M⁺ and M²⁺ cations, giving hydrides, oxides, hydroxides, carbonates, halides, and nitrides with predictable stoichiometry.

Most-tested idea: the diagonal relationship — Li resembles Mg, and Be resembles Al — because of comparable charge-to-size ratio (ρ = z/r).

High-yield pointers:

  • Hydration enthalpy ΔH_hyd ∝ −z²/r; lattice energy follows the Born–Landé relation.
  • Flame colours: Li crimson-red, Na yellow, K violet, Ca brick-red, Sr crimson, Ba apple green.
  • Group 2 sulphate solubility decreases down the group; hydroxide solubility increases down the group.

🟡 Standard — Regular Study (2d–2mo)

Standard content for students with a few days to months.

Electronic configuration and the +1/+2 oxidation state

Alkali metals adopt ns¹ and almost always exhibit the +1 oxidation state; alkaline earth metals adopt ns² and exhibit +2. The two outer electrons of Be/Mg are held tightly because of small atomic radii, so first and second ionisation enthalpies are higher in Group 2 than Group 1. Down either group, ionisation enthalpy decreases, atomic and ionic radii increase, and electronegativity falls.

Hydration enthalpy and lattice enthalpy

The two quantities compete in every solubility question:

ΔH_soln = ΔH_lattice + ΔH_hydration

A salt dissolves readily when the hydration enthalpy outweighs the lattice enthalpy. Charge density z/r controls both — small, highly charged ions attract water strongly but also pack tightly into the lattice.

Anomalous behaviour of Li and Be

Lithium differs from its heavier congeners because of its tiny radius and high charge density: it forms the covalent Li₂O and the nitride Li₃N directly with N₂, and Li₂CO₃ decomposes on heating (unlike Na₂CO₃). LiF and Li₂CO₃ are sparingly soluble. Beryllium shows analogous divergence — BeCl₂ is covalent (polymeric chains in solid, dimeric units in vapour), and Be(OH)₂ and BeO are amphoteric, dissolving in both acid and NaOH.

Group 2 saltTrend down the groupReason
Hydroxides M(OH)₂Solubility increases (Mg < Ca < Sr < Ba)Hydration gain dominates for small ions
Sulphates MSO₄Solubility decreases (Mg > Ca > Sr > Ba)Lattice enthalpy dominates for large ions
Carbonates MCO₃Solubility decreases; thermal stability also decreases down the groupHigher z/r polarises CO₃²⁻, easing decomposition

Exam pattern: expect 1–2 MCQs — a comparison table filler, anomalous behaviour identification, or flame-colour recall. Numerical questions are rare; conceptual reasoning over periodic trends is the default style.

🔴 Extended — Deep Study (3mo+)

Comprehensive coverage for students on a longer study timeline.

Diagonal relationship — charge density as the unifying principle

The diagonal pairings Li–Mg and Be–Al survive because their ρ = z/r values are nearly equal. Consequences:

  • Li and Mg both form nitrites on heating with N₂, give carbonates that decompose to the oxide + CO₂, and yield sparingly soluble fluorides and carbonates.
  • Be and Al both have amphoteric oxides/hydroxides, covalent chlorides that dimerise in vapour, and resist attack by water at room temperature.

Flame test and qualitative analysis

Flame colours arise from low-energy ns → (n−1)p electronic transitions whose wavelength shifts with the nuclear charge experienced by the valence electron. They are diagnostic for alkali metals (visible even in trace amounts) and for Ca, Sr, Ba. Sodium’s strong yellow emission must be masked with cobalt-blue glass when testing for potassium, whose violet line is otherwise drowned out.

Industrial and biological roles

Sodium is extracted by the Downs process (electrolysis of fused NaCl with CaCl₂ to lower the melting point) and converted to NaOH via the Castner–Kellner cell (mercury cathode, flowing Hg amalgam). Na₂CO₃ is produced by the Solvay process using NaCl, NH₃, and CaCO₃; the by-product CaCl₂ is the chief waste stream. Biologically, Na⁺/K⁺ gradients power nerve impulses and the Na⁺-K⁺-ATPase pump, while Ca²⁺ builds hydroxyapatite Ca₁₀(PO₄)₆(OH)₂ in bone/teeth and triggers actin–myosin coupling in muscle.

Common examiner traps

  • Decoy trend reversal: students assume all Group 2 salts become less soluble down the group; in reality hydroxides invert this rule.
  • Thermal stability ≠ solubility: SrCO₃ decomposes more easily than CaCO₃ because Sr²⁺ polarises CO₃²⁻ more weakly — a question often pairs these two ideas incorrectly.
  • BeCl₃ does not exist: writing BeCl₃ or assuming Be²⁺ behaves like the rest of Group 2 is a frequent error.

Worked micro-example: Compare ΔH_hyd of Li⁺ and Na⁺. With radii ≈ 76 pm and 102 pm, ΔH_hyd(Li⁺) ≈ −519 kJ/mol versus ΔH_hyd(Na⁺) ≈ −406 kJ/mol. The higher hydration enthalpy of Li⁺ is precisely why its salts deviate from the usual alkali-metal solubility order.

Practice prompts:

  1. Arrange Be(OH)₂, Mg(OH)₂, Ca(OH)₂, Sr(OH)₂, Ba(OH)₂ in order of increasing solubility, and justify using charge density.
  2. Explain why BeCl₂ is a covalent polymer while MgCl₂, CaCl₂, SrCl₂, and BaCl₂ are ionic.

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