Statement A — TRUE ✓
Lithium imparts crimson red colour to the flame. Flame test colours (NEET must-memorise): Li = crimson red, Na = golden yellow, K = violet/lilac, Rb = red-violet, Cs = blue. Li gives red because its electrons excite to specific energy levels and emit red-wavelength photons on returning.
Statement B — TRUE ✓
Sodium imparts golden yellow colour to flame — one of the most characteristic and intense flame colours. The Na D-line at 589 nm (yellow) is so bright that even trace amounts of Na contamination can mask other colours. This is used in street lights (sodium vapour lamps).
Statement C — FALSE ✗
Only Na, K, Rb, Cs are stored in kerosene oil (they react vigorously with air and moisture). Lithium is NOT stored in kerosene — it is less dense than kerosene and would float on it! Li is stored under dry mineral oil or in sealed containers under argon. Li reacts less vigorously than other alkali metals, so kerosene is unnecessary.
Statement D — TRUE ✓
Despite having the smallest size and highest ionisation energy, lithium is the strongest reducing agent among alkali metals (most negative E° = −3.04 V). This is due to its exceptionally high hydration enthalpy — Li⁺ is heavily hydrated in solution, releasing so much energy that it more than compensates for the high IE. This is an anomalous property of Li.
Alkali metals (Group 1: Li, Na, K, Rb, Cs, Fr) have one valence electron (ns¹) that they lose easily to form +1 ions. They are the most reactive metals, with reactivity increasing down the group. They are soft (can be cut with a knife), have low melting points (Cs melts at 28.5°C), low densities (Li, Na, K are less dense than water), and are excellent conductors of heat and electricity.
📌 Flame colours: Li = crimson red, Na = golden yellow, K = violet/lilac
📌 Rb = red-violet, Cs = blue
📌 Storage: Li in mineral oil/argon; Na, K in kerosene; Rb, Cs in sealed ampoules
📌 Density increases Li < Na < K (then K > Rb anomaly)
📌 All react with water: 2M + 2H₂O → 2MOH + H₂ (vigour increases Li → Cs)
Lithium resembles magnesium (diagonal relationship) more than Na/K due to similar charge/radius ratio. Anomalous properties of Li compared to other alkali metals:
📌 Highest ionisation energy, hardness, melting point, and density among alkali metals
📌 Strongest reducing agent despite highest IE — due to highest hydration enthalpy
📌 Li forms normal oxide (Li₂O), not peroxide or superoxide like Na, K
📌 Li reacts with N₂ (forms Li₃N) — other alkali metals don't react with N₂
📌 LiCl is covalent and deliquescent; NaCl/KCl are ionic
📌 Li₂CO₃ is unstable to heat (decomposes); Na₂CO₃/K₂CO₃ are stable
📌 Li is not stored in kerosene (too light — floats; also less reactive)
Reducing power in solution depends on: (1) Ionisation enthalpy (energy to remove electron from gas phase metal) — high for Li. (2) Sublimation enthalpy (solid → gas) — high for Li. (3) Hydration enthalpy (gaseous ion → aqueous ion) — highest for Li⁺ due to smallest size. The hydration enthalpy of Li⁺ (−519 kJ/mol) is so large that it more than compensates for the high IE and sublimation energy, giving Li the most negative standard electrode potential (E° = −3.04 V) among alkali metals. Hence Li is the strongest reducing agent.
2Na + 2H₂O → 2NaOH + H₂
Reactivity with water: Li < Na < K < Rb < Cs
Li: slow, steady; Na: fast, melts into ball; K: catches fire (violet flame)
Rb/Cs: explosive even with cold water
📌 NaOH (caustic soda): made by chlor-alkali process, used in soap, paper, textiles
📌 Na₂CO₃ (washing soda, Na₂CO₃·10H₂O): made by Solvay process, used as water softener, glass making
📌 NaHCO₃ (baking soda): decompose on heating → Na₂CO₃ + H₂O + CO₂; used in baking, antacids
📌 NaCl (common salt): mined or evaporated from sea water; raw material for Na, NaOH, Na₂CO₃, Cl₂, HCl
📌 Na₂O₂ (sodium peroxide): bleaching agent, used in submarines to regenerate O₂: 2Na₂O₂ + 2H₂O → 4NaOH + O₂
Alkaline earth metals (Be, Mg, Ca, Sr, Ba, Ra) have two valence electrons (ns²). They are harder, denser, and less reactive than alkali metals. They form +2 ions. Reactivity increases down the group — Be is least reactive, Ba most reactive. Be has anomalous properties and resembles Al (diagonal relationship). Flame colours: Ca = brick red, Sr = crimson, Ba = apple green.
📌 Flame colours: Ca = brick red, Sr = crimson red, Ba = apple/pale green
📌 Be and Mg don't impart any characteristic flame colour
📌 Be is amphoteric — dissolves in both acids and alkalis (like Al)
📌 Mg burns brilliantly in air — used in flares, incendiary bombs
📌 CaO (quicklime) + H₂O → Ca(OH)₂ (slaked lime, highly exothermic)
📌 CaCO₃ (limestone) → CaO + CO₂ at 1000°C (calcination)
Calcium compounds are extensively tested in NEET. Calcium oxide (CaO, quicklime) is obtained by heating limestone: CaCO₃ → CaO + CO₂. It reacts with water violently: CaO + H₂O → Ca(OH)₂ + heat. Calcium hydroxide (Ca(OH)₂, slaked lime) is used in whitewash, treating acidic soil, making mortar. Plaster of Paris (CaSO₄·½H₂O) sets hard: CaSO₄·½H₂O + 1½H₂O → CaSO₄·2H₂O (gypsum). Used for plastering walls, making casts for broken bones.
The Solvay (ammonia-soda) process makes sodium carbonate (washing soda) by: (1) Saturating NaCl solution with NH₃. (2) Passing CO₂: NaCl + NH₃ + CO₂ + H₂O → NaHCO₃↓ + NH₄Cl. (3) Filtering NaHCO₃, heating: 2NaHCO₃ → Na₂CO₃ + H₂O + CO₂. (4) Recovering NH₃: 2NH₄Cl + Ca(OH)₂ → CaCl₂ + 2NH₃ + 2H₂O. The only waste product is CaCl₂. NH₃ and CO₂ are recycled. The process is economical and produces very pure Na₂CO₃.