s-Block Elements Class 11 Notes

• Group 1 (alkali metals): Li, Na, K, Rb, Cs, Fr.
• Group 2 (alkaline earth metals): Be, Mg, Ca, Sr, Ba, Ra.
• They are strongly metallic, electropositive and form ionic compounds predominantly.

• One ns¹ electron; low ionisation enthalpy that decreases down the group (Cs has the least).
• Large atomic and ionic radii that increase down the group; the alkali metal ion is the smallest per electron shell, so it is heavily hydrated in water.
• Soft, low melting points because metallic bonding is weak with only one electron per atom.
• Flame colour: Li crimson, Na golden yellow, K lilac, Rb and Cs violet — from easily excited outer electrons.
• Cs is the most reactive; Li is the least.

• 2M + 2H₂O → 2MOH + H₂: reactivity with water increases down the group; Na reacts vigorously, K and Rb catch fire, Cs is explosive.
• They react with halogens to form ionic halides MX.
• They dissolve in liquid ammonia giving deep blue solutions that contain solvated electrons and are powerful reducing agents.
• They form ionic hydrides (LiH, NaH) with hydrogen — an exception to the usual molecular hydrides.

• Two ns² electrons; ionisation enthalpies higher than the alkali metals and decreasing down the group.
• M²⁺ ions are smaller and more heavily hydrated than the corresponding alkali metal ions; hydration enthalpy decreases from Be²⁺ to Ba²⁺.
• Harder, denser and higher melting than alkali metals because two electrons take part in metallic bonding.
• Less reactive than alkali metals; Mg and Be are relatively slow with water, Ca onwards react readily.
• Flame colour: Ca brick-red, Sr crimson, Ba apple-green; Be and Mg give no characteristic flame colour.

• Li (like Mg): forms a nitride Li₃N, carbonate that decomposes on heating, and a less soluble carbonate and hydroxide; both are harder than other group members.
• Be (like Al): forms covalent compounds, amphoteric oxide BeO, and stable complexes.
• Li and Be compounds are predominantly covalent, unlike the ionic compounds of the rest of their groups.

• Caustic soda NaOH: strong base, made by electrolysis of brine; soap, paper, textiles and bauxite extraction.
• Quicklime CaO (from limestone calcination) + water → slaked lime Ca(OH)₂.
• Plaster of Paris is hemihydrate 2CaSO₄·H₂O, made by heating gypsum CaSO₄·2H₂O at ~120 °C.

• Cement is mostly a mixture of calcium silicates and aluminates from limestone + clay.
• The key oxide components are CaO (~50%), SiO₂ and Al₂O₃.
• Cement sets when mixed with water: hydrates interlock and give strength over days; gypsum is added to regulate the setting time.

Example: Which alkali metal has the lowest ionisation enthalpy and why?

Solution: Caesium (Cs). Down the group the atomic radius grows and inner electrons shield the outer electron better, so the single ns¹ electron is pulled away more easily — ionisation enthalpy falls from Li to Cs.

Example: Why is Li⁺ the least mobile alkali metal ion in aqueous solution?

Solution: Li⁺ has the smallest ionic radius and the highest charge density, so it is hydrated the most. The large hydrated ion moves slowly, giving Li⁺ the lowest ionic mobility of the alkali metals.

Exercises (NCERT Q10.1–Q10.32)

• 10.1) What are the common physical and chemical features of alkali metals?
• 10.2) Discuss the general characteristics and gradation in properties of alkaline earth metals.
• 10.3) Why are alkali metals not found in nature?
• 10.4) Find out the oxidation state of sodium in Na₂O₂.
• 10.5) Explain why is sodium less reactive than potassium.
• 10.6) Compare the alkali metals and alkaline earth metals with respect to (i) ionisation enthalpy (ii) basicity of oxides and (iii) solubility of hydroxides.
• 10.7) In what ways lithium shows similarities to magnesium in its chemical behaviour?
• 10.8) Explain why can alkali and alkaline earth metals not be obtained by chemical reduction methods?
• 10.9) Why are potassium and caesium, rather than lithium used in photoelectric cells?
• 10.10) When an alkali metal dissolves in liquid ammonia the solution can acquire different colours. Explain the reasons for this type of colour change.
• 10.11) Beryllium and magnesium do not give colour to flame whereas other alkaline earth metals do so. Why?
• 10.12) Discuss the various reactions that occur in the Solvay process.
• 10.13) Potassium carbonate cannot be prepared by Solvay process. Why?
• 10.14) Why is Li₂CO₃ decomposed at a lower temperature whereas Na₂CO₃ at higher temperature?
• 10.15) Compare the solubility and thermal stability of the following compounds of the alkali metals with those of the alkaline earth metals. (a) Nitrates (b) Carbonates (c) Sulphates.
• 10.16) Starting with sodium chloride how would you proceed to prepare (i) sodium metal (ii) sodium hydroxide (iii) sodium peroxide (iv) sodium carbonate?
• 10.17) What happens when (i) magnesium is burnt in air (ii) quick lime is heated with silica (iii) chlorine reacts with slaked lime (iv) calcium nitrate is heated?
• 10.18) Describe two important uses of each of the following: (i) caustic soda (ii) sodium carbonate (iii) quicklime.
• 10.19) Draw the structure of (i) BeCl₂ (vapour) (ii) BeCl₂ (solid).
• 10.20) The hydroxides and carbonates of sodium and potassium are easily soluble in water while the corresponding salts of magnesium and calcium are sparingly soluble in water. Explain.
• 10.21) Describe the importance of the following: (i) limestone (ii) cement (iii) plaster of paris.
• 10.22) Why are lithium salts commonly hydrated and those of the other alkali ions usually anhydrous?
• 10.23) Why is LiF almost insoluble in water whereas LiCl soluble not only in water but also in acetone?
• 10.24) Explain the significance of sodium, potassium, magnesium and calcium in biological fluids.
• 10.25) What happens when (i) sodium metal is dropped in water? (ii) sodium metal is heated in free supply of air? (iii) sodium peroxide dissolves in water?
• 10.26) Comment on each of the following observations: (a) The mobilities of the alkali metal ions in aqueous solution are Li⁺ < Na⁺ < K⁺ < Rb⁺ < Cs⁺ (b) Lithium is the only alkali metal to form a nitride directly. (c) E⁰ for M²⁺(aq) + 2e⁻ → M(s) (where M = Ca, Sr or Ba) is nearly constant.
• 10.27) State as to why (a) a solution of Na₂CO₃ is alkaline? (b) alkali metals are prepared by electrolysis of their fused chlorides? (c) sodium is found to be more useful than potassium?
• 10.28) Write balanced equations for reactions between (a) Na₂O₂ and water (b) KO₂ and water (c) Na₂O and CO₂.
• 10.29) How would you explain the following observations? (i) BeO is almost insoluble but BeSO₄ is soluble in water (ii) BaO is soluble but BaSO₄ is insoluble in water (iii) LiI is more soluble than KI in ethanol.
• 10.30) Which of the alkali metal is having least melting point? (a) Na (b) K (c) Rb (d) Cs
• 10.31) Which one of the following alkali metals gives hydrated salts? (a) Li (b) Na (c) K (d) Cs
• 10.32) Which one of the alkaline earth metal carbonates is thermally the most stable? (a) MgCO₃ (b) CaCO₃ (c) SrCO₃ (d) BaCO₃

Answers at a glance

1. 1
   Q10.1 → Physical: soft, silvery-white, lustrous, low density and melting point, good conductors, low ionisation enthalpy. Chemical: strongly electropositive with a single ns¹ valence electron, always +1, highly reactive — form hydrides, oxides, hydroxides and salts with hydrogen, oxygen, halogens and water.
2. 2
   Q10.2 → Silvery, harder than alkali metals (metallic bonding), ns² and oxidation state +2; down the group: atomic/ionic radii increase while ionisation enthalpy, melting point and hydration energy decrease, and reactivity (and basicity of oxides/hydroxides) increases from Be to Ba.
3. 3
   Q10.3 → Because they are extremely reactive and reducing — they combine readily with the oxygen, water vapour and CO₂ of the air, so they exist only in combined state (as salts) and never free.
4. 4
   Q10.4 → +1. Na₂O₂ is a peroxide with the O₂²⁻ ion (each O is −1); for the anion to be −2, each sodium is +1.
5. 5
   Q10.5 → Potassium is bigger: its outer electron is farther from the nucleus and more shielded, so K has a much lower ionisation enthalpy than Na and loses its electron (and reacts) more easily.
6. 6
   Q10.6 → (i) Alkali metals have lower ionisation enthalpies (M → M⁺ is easier than M → M²⁺). (ii) Oxides of both are basic but alkali metal oxides are the more basic (e.g. K₂O more basic than CaO). (iii) Alkali metal hydroxides are freely soluble; alkaline earth hydroxides are only sparingly soluble (Be, Mg least; solubility increases down the group).
7. 7
   Q10.7 → Similar ionic/atomic size and charge density: both form chlorides soluble in alcohol and deliquescent (LiCl, MgCl₂), oxide and carbonate decompose on heating, both form nitrides directly (Li₃N, Mg₃N₂), give weak basic hydroxides, and polarise anions (covalent character, e.g. Li₂CO₃ like MgCO₃).
8. 8
   Q10.8 → They are themselves the strongest reducing agents known; no common reducing agent can donate electrons to their ions, so they are obtained only by electrolysis of their fused (molten) chlorides.
9. 9
   Q10.9 → Li has a relatively high ionisation enthalpy, so visible light cannot eject its electrons; K and Cs have very low ionisation enthalpies, so even visible light easily ejects electrons (photoelectric effect).
10. 10
    Q10.10 → Alkali metals ionise in liquid NH₃ giving ammoniated cations and ammoniated electrons, e⁻(am); the ammoniated electron absorbs light in the visible region — a dilute solution is blue, and a concentrated solution appears copper-bronze as ion pairs form.
11. 11
    Q10.11 → Be²⁺ and Mg²⁺ have very high ionisation enthalpies — the energy of the flame is too small to excite their electrons to higher levels — so they emit no visible colour; the heavier members are easily excited and emit their characteristic flame colours.
12. 12
    Q10.12 → NaCl + NH₃ + CO₂ + H₂O → NaHCO₃↓ + NH₄Cl (sodium bicarbonate precipitates); 2NaHCO₃ →Δ Na₂CO₃ + H₂O + CO₂ (calcination); the NH₃ is regenerated as 2NH₄Cl + Ca(OH)₂ → CaCl₂ + 2NH₃ + 2H₂O, and CO₂ is recycled.
13. 13
    Q10.13 → Because potassium bicarbonate KHCO₃ is much more soluble in water than NaHCO₃, so it does not precipitate out and cannot be separated by this process.
14. 14
    Q10.14 → Li⁺ is very small with a high charge density, so it strongly polarises the CO₃²⁻ ion, destabilising lithium carbonate (more covalent character) — hence it decomposes easily; Na₂CO₃ is more ionic and far more stable, decomposing only at very high temperature.
15. 15
    Q10.15 → (a) Alkali nitrates give nitrite + O₂ on heating, stable and soluble; alkaline earth nitrates decompose to oxide + NO₂ + O₂. (b) Alkali carbonates are soluble and thermally very stable; alkaline earth carbonates are sparingly soluble and decompose to MO + CO₂, stability increasing Be < Mg < Ca < Sr < Ba. (c) Alkali sulphates are soluble and stable; alkaline earth sulphates are sparingly soluble (solubility falling from BeSO₄ to BaSO₄) and decompose only at high temperature.
16. 16
    Q10.16 → (i) Electrolysis of fused NaCl (Down's cell): 2NaCl → 2Na + Cl₂. (ii) Electrolysis of brine: 2NaCl + 2H₂O → 2NaOH + Cl₂ + H₂. (iii) Burn sodium in excess air/oxygen: 2Na + O₂ → Na₂O₂. (iv) Solvay process (NaCl + NH₃ + CO₂ + H₂O → NaHCO₃ →Δ Na₂CO₃).
17. 17
    Q10.17 → (i) Mg forms MgO (with some Mg₃N₂): 2Mg + O₂ → 2MgO; 3Mg + N₂ → Mg₃N₂. (ii) CaO + SiO₂ → CaSiO₃ (calcium silicate). (iii) Ca(OH)₂ + Cl₂ → CaOCl₂ + H₂O (bleaching powder). (iv) 2Ca(NO₃)₂ →Δ 2CaO + 4NO₂ + O₂.
18. 18
    Q10.18 → (i) Caustic soda: manufacture of soap and detergents, and of paper and rayon (also petroleum refining). (ii) Sodium carbonate: manufacture of glass, and softening of hard water (also soap and washing soda). (iii) Quicklime: manufacture of cement and as the basic flux in iron/steel making (also water softening and soil treatment).
19. 19
    Q10.19 → (i) Vapour: a linear monomer, Cl–Be–Cl, with sp-hybridised Be. (ii) Solid: a polymeric chain in which each Be is tetrahedrally surrounded (sp³) by four chlorine atoms forming bridging bonds.
20. 20
    Q10.20 → For Na/K salts the hydration energy released outweighs the lattice energy, so dissolution is favourable; for Mg/Ca salts the much higher lattice energy is not compensated by hydration energy, so they stay sparingly soluble.
21. 21
    Q10.21 → (i) Limestone: raw material for lime, cement, and the flux in iron/steel and glass manufacture. (ii) Cement: the binding material of buildings, bridges and roads (limestone + clay). (iii) Plaster of paris (CaSO₄·½H₂O): setting broken bones, making casts, moulds, and false ceilings.
22. 22
    Q10.22 → Li⁺ is the smallest, so it has the greatest charge density and its hydration energy is the largest — it traps water molecules, giving hydrated salts like LiCl·H₂O; the larger alkali ions hydrate weakly and form anhydrous salts.
23. 23
    Q10.23 → LiF has very high lattice energy (both ions small) that hydration cannot overcome, so it is insoluble; LiCl's lattice is loose and its small Li⁺ hydrates strongly, and the covalent character of LiCl makes it dissolve in acetone too.
24. 24
    Q10.24 → Na⁺ and K⁺ maintain osmotic balance and the Na⁺/K⁺ pump that transmits nerve impulses; Ca²⁺ is needed for blood clotting, muscle contraction and membrane/nerve function; Mg²⁺ activates enzymes and is the central ion of chlorophyll.
25. 25
    Q10.25 → (i) Violent reaction: 2Na + 2H₂O → 2NaOH + H₂. (ii) Sodium burns to the peroxide: 2Na + O₂ → Na₂O₂. (iii) Na₂O₂ + 2H₂O → 2NaOH + H₂O₂ (which further decomposes to O₂).
26. 26
    Q10.26 → (a) The aqueous (hydrated) ionic radius is largest for Li⁺ and smallest for Cs⁺, so mobility increases down the group. (b) Li⁺ is small enough to stabilise the highly charged N³⁻ ion in a lattice (high lattice energy of Li₃N) — the larger alkali metals cannot. (c) The rise in ionisation energy down the group is almost exactly offset by the falling hydration energy of M²⁺, so E⁰ stays nearly constant (~ −2.9 V).
27. 27
    Q10.27 → (a) CO₃²⁻ hydrolyses: CO₃²⁻ + H₂O ⇌ HCO₃⁻ + OH⁻. (b) Their ions are too strong a reducing action for any chemical reduction and their aqueous solutions would reduce water; only electrolysis of the fused chloride works. (c) Sodium is abundant, cheaper and less reactive/more manageable than potassium.
28. 28
    Q10.28 → (a) Na₂O₂ + 2H₂O → 2NaOH + H₂O₂ (b) 2KO₂ + 2H₂O → 2KOH + H₂O₂ + O₂ (c) Na₂O + CO₂ → Na₂CO₃.
29. 29
    Q10.29 → (i) BeO has very high lattice energy that hydration cannot beat; BeSO₄'s sulphate anion is large, lowering lattice energy, so hydration wins. (ii) With BaO the high hydration energy of the small oxide ion overcomes the lattice, but in BaSO₄ both ions are large and the lattice energy is huge → insoluble. (iii) Li⁺ is small and I⁻ large, so LiI has pronounced covalent character (Fajans' rule) and dissolves in the organic solvent ethanol; KI is ionic and does not.
30. 30
    Q10.30 → (d) Cs ~ 28.5 °C — melting point of alkali metals decreases down the group as atoms grow and metallic bonding weakens.
31. 31
    Q10.31 → (a) Li — its small size gives the highest hydration energy, so lithium salts crystallise with water (e.g. LiCl·H₂O).
32. 32
    Q10.32 → (d) BaCO₃ — thermal stability of the alkaline earth carbonates increases down the group (BeCO₃ decomposes most easily).

MCQ Questions

1 Mark Questions

1. 1
   Why are alkali metals not found in free nature? (Ans: they are extremely reactive/reducing, so they combine with O₂, water and CO₂ and exist only in combined state)
2. 2
   Give the general electronic configuration of alkali metals and alkaline earth metals. (Ans: ns¹ and ns²)
3. 3
   Why are K and Cs, rather than Li, used in photoelectric cells? (Ans: their low ionisation enthalpy lets visible light eject electrons easily)
4. 4
   What is the Li–Mg diagonal relationship? (Ans: Li resembles Mg because of similar size and charge density, giving similar chemical behaviour)

2 Mark Questions

1. 1
   Why is LiF almost insoluble in water while LiCl is soluble in water and acetone? (Ans: LiF has high lattice energy; LiCl has covalent character and strong hydration of Li⁺)
2. 2
   Why do Be and Mg not give flame colour? (Ans: their high ionisation enthalpy cannot be overcome by flame energy, so no excitation/emission)
3. 3
   Why is Li₂CO₃ decomposed at a lower temperature than Na₂CO₃? (Ans: Li⁺ has high polarising power so it is more covalent/less stable; Na₂CO₃ is ionic and very stable)
4. 4
   Explain why a solution of Na₂CO₃ is alkaline. (Ans: CO₃²⁻ hydrolyses: CO₃²⁻ + H₂O ⇌ HCO₃⁻ + OH⁻)

3 Mark Questions

1. 1
   Discuss the various reactions involved in the Solvay process. (Ans: NaCl + NH₃ + CO₂ + H₂O → NaHCO₃↓; calcination: 2NaHCO₃ → Na₂CO₃ + CO₂ + H₂O; NH₃ regenerated via Ca(OH)₂)
2. 2
   Why cannot alkali and alkaline earth metals be obtained by chemical reduction? (Ans: they are the strongest reducing agents; their ions cannot be reduced — only electrolysis of fused chlorides works)
3. 3
   In what ways does Li show similarities to Mg? (Ans: both form nitrides directly, soluble chlorides in alcohol, decomposing carbonates, weak basic hydroxides — due to similar sizes)
4. 4
   Compare alkali metals and alkaline earth metals in respect of ionisation enthalpy, basicity of oxides and solubility of hydroxides. (Ans: lower IE for alkali, more basic oxides, soluble hydroxides; opposite trends for alkaline earths)

5 Mark Questions

1. 1
   Discuss the general characteristics and gradation in properties of alkaline earth metals from Be to Ba — atomic/ionic radii, ionisation enthalpy, melting point, reactivity and thermal stability of carbonates, with reasons.
2. 2
   Explain the Solvay process for the manufacture of sodium carbonate in detail — raw materials, all the reactions, and why potassium carbonate cannot be made by this method. (Ans: KHCO₃ is too soluble to precipitate)
3. 3
   Discuss the alkali metal–liquid ammonia solutions — why they are blue and become copper-bronze, their electrical conductivity and chemical reducing action — and the electrolytic methods for obtaining sodium, sodium hydroxide and sodium peroxide from sodium chloride.
4. 4
   Starting from sodium chloride, describe how sodium metal, sodium hydroxide, sodium peroxide and sodium carbonate are prepared, writing balanced equations and naming the process/electrode reactions for each.

Assignments walk the chapter's topics in order; the tests follow the current CBSE blueprint (MCQ/VSA, 2M, 3M and 5M questions) with answer keys. They will appear here as direct download links.