Group 1

Alkali Metals (Group 1: Li, Na, K, Rb, Cs; Fr—radioactive)

1) Introduction

• s-Block; valence ns¹, highly reactive; hydroxides (MOH) are strongly alkaline.

• (Alkali metals): Li, Na, K, Rb, Cs, Fr. Called alkali metals because they form strongly alkaline hydroxides with water.

2) Occurrence, Extraction, and Uses

Occurrence

• Li: spodumene LiAl(SiO₃)₂, lepidolite.

• Na: rock salt NaCl, Chile saltpetre NaNO₃, cryolite Na₃AlF₆.

• K: carnallite KCl·MgCl₂·6H₂O, saltpetre KNO₃, kainite KCl·MgSO₄·3H₂O.

• Rb/Cs: traces with Li/K minerals. Fr: trace from Ac α-decay (longest-lived isotope ²²³Fr, t½ ≈ 21 min).

Extraction (key routes)

• Li/Na: electrolysis of fused chlorides.

• K: displacement volatility process: Na(g)+K⁺(l) ⇌ Na⁺(l)+K(g); K distils off.

• Rb/Cs: reduction of chlorides with Ca at ~1000 K (vacuum).

Representative Uses (high-yield)

• Na (liquid): reactor coolant; Na vapor lamps.

• Cs: photoelectric cells (low work function).

• Li: stearate greases; LiAlH₄ (reducing agent), batteries.

• NaBH₄ (reducing agent).

• NaCl → NaOH/Na₂CO₃/Cl₂/H₂; KOH (liquid detergents), KO₂ (rebreathers: CO₂ scrubbing + O₂ release), KClO₃ (matches), KNO₃ (gunpowder), K-fertilizers.

3) General Characteristics

3.1 Physical Trends (down the group)

• Size (atomic/ionic): ↑ (largest in their periods); M⁺ < M.

• Density: generally ↑; exception: K < Na.

• m.p./b.p. & hardness: ↓ (one loosely held valence electron).

• Conductivity (thermal/electrical): high.

• IE₁ & electronegativity: very low; IE₁ ↓ down group → reactivity ↑.

Flame colours (nm): Li crimson 670.8; Na yellow 589.2; K violet 766.5; Rb red-violet 780.0; Cs blue 455.5.

Standard potentials E°(M⁺/M) ~: Li −3.04 > Na −2.71 > K −2.925 ≈ Rb −2.93 ≈ Cs −2.927 (very strong reducers; Li strongest overall due to hydration).

3.2 Solubility Logic (Lattice vs Hydration Energy)

• Small anions (F⁻, O²⁻, OH⁻): lattice energy falls sharply down group → solubility ↑ (e.g., LiF < … < CsF).

• Large anions (I⁻, NO₃⁻, SO4--): hydration drop dominates → solubility ↓ down group (often Li⁺ salts most soluble here).

• Size-match rule: small–small / large–large stabilize lattice; mismatch favors solubility.

4) Chemical Properties and Key Compounds

4.1 Oxides / Peroxides / Superoxides (air burn)

• Li → Li₂O (some Li₂O₂); Na → Na₂O₂ (some NaO₂); K/Rb/Cs → MO₂.

• Stability of O₂²⁻/O₂⁻ ↑ down group (large anions stabilized by larger cations).

• Hydrolysis:
  Li₂O + H₂O → 2 LiOH
  Na₂O₂ + 2 H₂O → 2 NaOH + H₂O₂
  2 KO₂ + 2 H₂O → 2 KOH + H₂O₂ + O₂
  (H₂O₂ formation → oxidizing behavior in water)

• Superoxides: yellow–orange; paramagnetic (O₂⁻ has 1 unpaired e⁻). Na₂O₂ widely used as an oxidant.

4.2 Hydroxides

• Strongest bases; basic strength ↑ down group (lower M⁺ charge density).

4.3 Hydrides (ionic, MH)

• 2 M + H₂ → 2 MH (≈673 K; Li ~1073 K).

• MH + H₂O → MOH + H₂↑.

• Synthesis links: 4 LiH + AlCl₃ → LiAlH₄ + 3 LiCl; NaBH₄ via borate routes.

4.4 Sulphides & Polysulphides

• M₂S and M₂Sₙ (n=2–6); zig-zag S chains.

4.5 Carbides/Acetylides

• Li₂C₂ from C or C₂H₂; M₂C₂ + 2 H₂O → 2 MOH + C₂H₂↑.

4.6 Reactivity Snapshots

• With O₂/air: tarnish; burn vigorously; Li forms Li₃N uniquely → store under kerosene.

• With water: 2 M + 2 H₂O → 2 M⁺ + 2 OH⁻ + H₂↑ (Li slower despite most-negative E° due to high hydration energy).

• With X₂: MX (LiX more covalent; LiI most).

• With ROH: M + ROH → MOR + ½ H₂ (stronger down group).

• Liquid NH₃: deep blue, conducting, paramagnetic (solvated e⁻); on standing → amide + H₂; concentrated → bronze & diamagnetic.

• Only Li readily forms nitride Li₃N at r.t.

5) Thermal Stability & Decomposition Patterns

• Small-anion salts (MF, M₂O, MOH): stability ↓ down group (lattice energy drop).

• Large-anion salts (I⁻, NO₃⁻): stability ↑ down group.

• Carbonates: Li₂CO₃ decomposes (→ Li₂O + CO₂); others resist Bunsen flame.

• Bicarbonates: all exist as solids except LiHCO₃ (only in solution).

• Hydroxides: LiOH partially unstable on strong heating (→ Li₂O); others stable.

• Nitrates: LiNO₃ → Li₂O + NO₂ + O₂; others → MNO₂ + O₂.

6) Solvation/Hydration & Mobility

• Degree of hydration: Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺.

• Hydrated radius: Li⁺(aq) largest → lowest ionic mobility; conductance in water: Cs⁺ > Rb⁺ > K⁺ > Na⁺ > Li⁺.
  Mnemonic: “Large dry, tiny wet—Li⁺ is tiniest dry but largest when hydrated.”

7) Complexation Behaviour

• Alkali metals are poor complexers (large, +1, low charge density). Stability trend: Li > Na > K > Rb > Cs.

• Complexes mainly with chelates (β-diketones, nitrophenols, nitrosonaphthols).

• Li⁺ forms more adducts: e.g., [Li(NH₃)₄]⁺, [LiCl(py)₂(ROH)].

8) Anomalous Behaviour of Lithium (and Li–Mg similarity)

8.1 Why anomalous?

• Very small size & high polarising power (charge/radius) → more covalent compounds; diagonal relationship with Mg²⁺.

8.2 Differences: Li vs other alkalis

• Harder; higher m.p./b.p.

• Least reactive in many media but strongest reducer in water (very high hydration enthalpy).

• Burns to Li₂O; forms Li₃N directly; LiF, Li₂O less soluble than Na/K analogues.

• LiCl deliquescent; forms LiCl·2H₂O; no solid LiHCO₃; no Li ethynide; LiNO₃ → Li₂O (others → nitrites).

8.3 Similarities: Li ↔ Mg (diagonal relationship)

• Both harder & lighter; react slowly with water.

• Oxides/hydroxides less soluble; hydroxides decompose on heating.

• Both form nitrides; no superoxides.

• Carbonates decompose on heating; no solid hydrogencarbonates.

• Chlorides (LiCl, MgCl₂) soluble in ethanol, deliquescent, crystallise as hydrates (LiCl·2H₂O, MgCl₂·8H₂O).

9) Important Sodium Compounds (prep → properties → uses)

9.1 Sodium carbonate (washing soda) — Na₂CO₃·10H₂O

• Solvay process:
  NH₃ + CO₂ + H₂O → NH₄HCO₃;
  NH₄HCO₃ + NaCl → NaHCO₃↓ + NH₄Cl;
  2 NaHCO₃ (Δ) → Na₂CO₃ + CO₂ + H₂O;
  NH₃ recovery: NH₄Cl + Ca(OH)₂ → NH₃ + CaCl₂ + H₂O.
  Not viable for K₂CO₃ (KHCO₃ too soluble).

• Hydrates: deca → mono (~375 K) → soda ash (>373 K).

• Alkaline solution: CO₃²⁻ + H₂O ⇌ HCO₃⁻ + OH⁻.

• Uses: water softening; glass/soap/borax/NaOH; paper, paints, textiles; lab reagent.

9.2 Sodium chloride — NaCl

• Source: sea water (2.7–2.9%); solar evaporation.

• Purification: dissolve crude salt, filter; saturate with HCl(g) → NaCl crystallises (CaCl₂/MgCl₂ remain).

• Uses: table salt; source for Na₂O₂, NaOH, Na₂CO₃.

9.3 Sodium hydroxide (caustic soda) — NaOH

• Castner–Kellner cell (Hg cathode):
  Na⁺ + e⁻ → Na(Hg); at anode: Cl₂↑;
  Amalgam + H₂O → NaOH + H₂ + Hg(recycled).

• Properties: white, deliquescent; absorbs CO₂ → Na₂CO₃.

• Uses: soaps, paper, rayon; petroleum refining; bauxite purification; mercerising cotton; fats/oils; lab reagent.

9.4 Sodium hydrogencarbonate (baking soda) — NaHCO₃

• Prep: saturate Na₂CO₃(aq) with CO₂ → NaHCO₃↓.

• Uses: baking (CO₂ release on heating), mild antiseptic, fire extinguishers.

10) Biological Importance of Na⁺ and K⁺

• In a 70-kg adult: ~90 g Na, 170 g K.

• Na⁺: mainly extracellular (plasma/interstitial fluid) → nerve transmission, osmotic balance, glucose/AA co-transport.

• K⁺: mainly intracellular → enzyme activation, glucose oxidation → ATP, nerve transmission (with Na⁺).

• Typical gradients (mmol L⁻¹): Plasma: Na⁺ ~143, K⁺ ~5; RBC interior: Na⁺ ~10, K⁺ ~105.

• Na⁺/K⁺ pump: maintains gradients; >⅓ of resting ATP consumption (≈ 15 kg ATP/day in a resting human).