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🧪 Chemistry · Class 11 · NEET & JEE

s-Block Elements

Group 1 (alkali metals) and Group 2 (alkaline earth metals). Learn their reactions with water and air, important compounds like NaOH, Na₂CO₃, and CaCO₃, and anomalous behaviour of lithium and beryllium.

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Reading time~8 min

Revision time~3 min

Last updated2026-07-19

1 Read the chapter ~8 min

🎯 Key Points

Group 1 (alkali metals): 1 valence electron, +1 ion, softest metals, react vigorously with water
Group 2 (alkaline earth metals): 2 valence electrons, +2 ion, less reactive than Group 1
Li is anomalous (resembles Mg); Be is anomalous (resembles Al) — both diagonal relationships
Thermal stability of carbonates/nitrates increases down the group; hydroxide solubility increases down Group 2; sulphate solubility decreases down Group 2
Key compounds: NaOH (caustic soda), Na₂CO₃ (washing soda), NaHCO₃ (baking soda), CaO (quicklime), Ca(OH)₂ (slaked lime), Plaster of Paris
Flame test colours: Li=crimson, Na=golden yellow, K=violet, Ca=brick red, Ba=apple green

Each s-block metal ion emits a characteristic flame colour because heating excites electrons to higher energy levels; light is emitted as they fall back, with the colour determined by the specific energy gap.

Group 1: Alkali Metals (Li, Na, K, Rb, Cs)

Softest metals with the lowest ionisation energies. Highly reactive due to their single valence electron.

React vigorously with water: 2Na + 2H₂O → 2NaOH + H₂ (exothermic)
Na and K stored under kerosene to prevent reaction with air
Reactivity increases down the group: Li < Na < K < Rb < Cs
Flame colours: Li = crimson red, Na = golden yellow, K = violet

Group 2: Alkaline Earth Metals (Be, Mg, Ca, Sr, Ba)

Two valence electrons; form +2 ions; less reactive than Group 1
Be and Mg do not react with cold water; Ca and below react readily
Reactivity increases down the group

Important Compounds

NaOH (caustic soda): Strong base; used in soap, paper, textile industries
Na₂CO₃ (washing soda): Na₂CO₃.10H₂O; used in glass, detergents, water softening
NaHCO₃ (baking soda): Used in baking, fire extinguishers, antacids
CaO (quicklime): Formed by heating CaCO₃; reacts with water to form slaked lime
Ca(OH)₂ (slaked lime): Used in construction and whitewash
Plaster of Paris: CaSO₄.½H₂O; sets hard by absorbing water

Anomalous Behaviour

Li resembles Mg (diagonal relationship in periodic table)
Be resembles Al (diagonal relationship)
Small size and high charge density cause differences from rest of group

Anomalous Behaviour of Lithium

Smallest size and highest charge density among alkali metals, giving it unusually strong polarising power
Li is the only alkali metal that reacts directly with N₂ to form a nitride (Li₃N); other alkali metals do not
LiCl is covalent and soluble in organic solvents, unlike the ionic chlorides of other alkali metals
LiHCO₃ does not exist as a solid (unlike NaHCO₃), and Li salts are more often hydrated due to high hydration enthalpy

Anomalous Behaviour of Beryllium

Be salts are covalent in nature (not ionic) because of its very small size and high polarising power
BeO is amphoteric, while oxides of other Group 2 elements are basic
Be does not show coordination number 6 in its compounds (too small), unlike Mg and Ca
Be and its compounds are toxic, unlike other alkaline earth elements

Solubility and Thermal Stability Trends

Solubility of hydroxides increases down Group 2: Be(OH)₂ < Mg(OH)₂ < Ca(OH)₂ < Sr(OH)₂ < Ba(OH)₂
Thermal stability of carbonates and nitrates increases down both groups, since larger cations stabilise the large carbonate/nitrate anion better (smaller cations polarise and distort the anion, making it decompose more easily on heating)
Solubility of sulphates decreases down Group 2: BeSO₄ and MgSO₄ are soluble, BaSO₄ is insoluble (used in the gravimetric test for SO₄²⁻ and as a contrast medium in X-rays)

Biological Importance

Na⁺ and K⁺ maintain osmotic balance, nerve impulse transmission, and ionic balance in cells
Mg²⁺ is the central atom in chlorophyll, essential for photosynthesis
Ca²⁺ is essential for bone and teeth formation, blood clotting, and muscle contraction

General Electronic Configuration and Periodic Trends

Configuration: Group 1 = [noble gas]ns¹; Group 2 = [noble gas]ns²
Atomic and ionic radii: the largest in their respective periods and increase down each group as new shells are added
Ionisation enthalpy: lowest in each period and decreases down the group (Group 2 > Group 1 because of higher nuclear charge and smaller size); the second IE of Group 2 metals is still low enough for them to form +2 ions
Hydration enthalpy: decreases with increasing ionic size, so Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺; this is why Li⁺ salts are the most heavily hydrated and why Li is the strongest reducing agent in aqueous solution despite its high IE
Metallic character: low melting points and low densities (Li, Na, K float on water) because weak metallic bonding results from just one or two loosely held valence electrons

Reactions with Air and Oxygen (Oxides, Peroxides, Superoxides)

Lithium forms mainly the normal oxide: 4Li + O₂ → 2Li₂O
Sodium forms the peroxide: 2Na + O₂ → Na₂O₂
Potassium, rubidium and caesium form the superoxide: K + O₂ → KO₂; the superoxide ion O₂⁻ is paramagnetic and a strong oxidiser
KO₂ is used in space capsules and submarines to absorb CO₂ and release O₂: 4KO₂ + 2CO₂ → 2K₂CO₃ + 3O₂
All s-block metals react with water liberating H₂ and forming hydroxides; reactivity toward water increases down each group

Solutions in Liquid Ammonia

Alkali (and the heavier alkaline earth) metals dissolve in liquid ammonia to give deep blue solutions that conduct electricity
The blue colour and conductivity are due to ammoniated (solvated) electrons, [e(NH₃)ₙ]⁻, released when the metal ionises in the solvent
Dilute solutions are paramagnetic and act as powerful reducing agents; concentrated solutions become bronze-coloured and diamagnetic

Diagonal Relationship (Li-Mg and Be-Al)

The first element of Groups 1 and 2 resembles the diagonally placed second-group element because of a similar charge/size ratio (charge density) and similar polarising power
Li and Mg: both react with N₂ to form nitrides (Li₃N, Mg₃N₂), both give normal oxides, both have carbonates that decompose on heating, and their chlorides are deliquescent and soluble in organic solvents
Be and Al: both form covalent, amphoteric oxides (BeO, Al₂O₃), both dissolve in alkali releasing H₂, and their chlorides (BeCl₂, AlCl₃) are covalent, Lewis-acidic and act as Friedel-Crafts catalysts

🚀 JEE Advanced Edge

Why flame colours occur: Heat excites a valence electron to a higher energy level; as it falls back, it emits a photon of energy exactly equal to the energy gap, corresponding to a specific visible wavelength/colour. Larger atoms (Cs, K) have more closely-spaced energy levels, often giving violet/UV-shifted colours, while smaller, more tightly bound electrons (Li) give higher-energy red emission.

Why LiCl is covalent (Fajan's rule application): Li⁺ is extremely small with very high charge density, so it strongly polarises the large Cl⁻ ion's electron cloud, pulling shared electron density toward itself and giving the bond significant covalent character — exactly the Fajan's rule scenario (small, highly-charged cation + large anion = more covalent).

Worked problem: Explain why BeCl₂ is a covalent, electron-deficient molecule that exists as a dimer in vapour phase. Approach: Be has only 2 valence electrons (2 bonds, no lone pairs) and a very small size; BeCl₂ in vapour phase has only 4 electrons around Be (electron deficient, like BF₃). To complete its octet, two BeCl₂ units share Cl atoms via coordinate bonds, forming a bridged dimer (Be₂Cl₄) where each Cl donates a lone pair to the adjacent Be.

2 Revise ~3 min before the exam

📐 Formula Sheet

Group 1 (alkali metals): ns¹, +1 ions, softest and most reactive metals; reactivity increases down the group
Group 2 (alkaline earth): ns², +2 ions, harder and less reactive than group 1
Flame colours: Li crimson, Na yellow, K lilac; Ca brick-red, Sr crimson, Ba apple-green
Reactivity with water: increases down each group; Li reacts gently, Cs violently
Solubility trends: group 2 hydroxides become more soluble down the group; sulfates become less soluble
Diagonal relationship: Li resembles Mg; Be resembles Al
Anomalies: Li and Be differ from their groups (small size, high polarising power)
Key compounds: washing soda Na₂CO₃·10H₂O, baking soda NaHCO₃, lime CaO, gypsum CaSO₄·2H₂O, plaster of Paris CaSO₄·½H₂O

3 Practice apply it

✍️ Worked Examples

Example 1 — Reactivity trend
Q: Why does reactivity with water increase from Li to Cs?
Step 1 — Reactivity depends on how easily the ns¹ electron is lost, i.e. the ionisation energy.
Step 2 — Going down, atoms get bigger and the outer electron is farther from the nucleus and better shielded.
Step 3 — Ionisation energy falls, so the electron is lost more easily and reactions get more vigorous.
Answer: lower ionisation energy down the group makes the metals more reactive. Note: caesium reacts explosively, lithium only steadily.

Example 2 — Diagonal relationship
Q: Give two ways lithium resembles magnesium rather than the other alkali metals.
Step 1 — Li and Mg have similar charge-to-size ratios (polarising power), placing them diagonally.
Step 2 — Both form nitrides directly with N₂: 6Li + N₂ → 2Li₃N, unlike Na or K.
Step 3 — Both have covalent, water-soluble-limited carbonates that decompose on heating, unlike the very stable Na₂CO₃.
Answer: lithium forms a nitride and has an unstable carbonate, like magnesium. Key idea: the diagonal relationship arises from matching polarising power.

Example 3 — Plaster of Paris
Q: Write the reaction for setting plaster of Paris and explain why it hardens.
Step 1 — Plaster of Paris is CaSO₄·½H₂O.
Step 2 — On adding water it rehydrates: CaSO₄·½H₂O + 1½H₂O → CaSO₄·2H₂O (gypsum).
Step 3 — The interlocking gypsum crystals set into a hard, slightly expanded solid.
Answer: it rehydrates to gypsum, whose crystal network gives it strength — which is why it is used in casts and moulds.