s-Block Elements
s-Block Elements for JEE Main & Advanced
Alkali Metals (Group 1)
General Trends and PropertiesTopic 1
Alkali Metals: Group 1 elements — Li, Na, K, Rb, Cs, Fr. Outer config: $ns^1$. Soft, low-melting, highly reactive metals.
Electronic Configuration:
| Element | $Z$ | Config |
|---|---|---|
| Li | 3 | $[He]\,2s^1$ |
| Na | 11 | $[Ne]\,3s^1$ |
| K | 19 | $[Ar]\,4s^1$ |
| Rb | 37 | $[Kr]\,5s^1$ |
| Cs | 55 | $[Xe]\,6s^1$ |
| Fr | 87 | $[Rn]\,7s^1$ (radioactive) |
Trends Down the Group:
| Property | Trend |
|---|---|
| Atomic radius | Increases ↑ |
| Ionic radius (M⁺) | Increases ↑ |
| Ionization enthalpy | Decreases ↓ |
| Electronegativity | Decreases ↓ |
| Density | Generally increases (K is exception; less than Na) |
| Melting/Boiling point | Decreases ↓ (weakening of metallic bond) |
| Reactivity with water | Increases (Li slow, Cs explosive) |
| Hydration enthalpy of M⁺ | Decreases (Li⁺ most hydrated) |
Physical Properties:
- All are soft; can be cut with a knife
- Silvery-white luster (freshly cut); tarnish in air
- Low density (Li < water; Na, K float on water)
- Low melting points (Cs MP = $28.5°$C, melts in hand!)
- Good conductors of heat and electricity
- Emit characteristic flame colors:
- Li: crimson red
- Na: golden yellow
- K: violet (lilac)
- Rb: red-violet
- Cs: blue
(Loose outer electron easily excited, emits visible light on relaxation.)
Chemical Properties:
- With water: $2M + 2H_2O \to 2MOH + H_2$ (vigorous, exothermic)
- With air: Tarnish quickly forming oxides; burn in O₂:
- $4Li + O_2 \to 2Li_2O$ (oxide)
- $2Na + O_2 \to Na_2O_2$ (peroxide)
- $K, Rb, Cs + O_2 \to MO_2$ (superoxide; gives KO₂, RbO₂, CsO₂)
- With halogens: $2M + X_2 \to 2MX$ (typically ionic salts)
- With hydrogen: $2M + H_2 \to 2MH$ (ionic hydrides)
- With ammonia (liquid): Dissolve to give deep-blue, paramagnetic, conducting solution containing solvated electron: $M + (x+y)NH_3 \to [M(NH_3)_x]^+ + [e(NH_3)_y]^-$. Color due to ammoniated electrons.
- Reducing power: Strong reducing agents; $Li$ is strongest (highest -ve $E^\ominus$, despite Cs having lowest IE — due to high hydration of Li⁺).
Why is Li the strongest reducing agent in Group 1 even though Cs has the lowest ionization enthalpy?
Show solution
Reducing power in solution depends on three energies: sublimation, ionization, hydration. While Cs has lowest IE, Li has the highest hydration enthalpy (small Li⁺ heavily hydrated). The very negative hydration energy of Li⁺ makes the overall process most favorable, hence most negative $E^\ominus$ → strongest reducing agent in aqueous solution.
Final Answer: High hydration enthalpy of small Li⁺ ion makes Li strongest RA.
Predict products: $2Na + O_2 \to ?$ vs $K + O_2 \to ?$
Show solution
$2Na + O_2 \to Na_2O_2$ (sodium peroxide) $K + O_2 \to KO_2$ (potassium superoxide) The larger cations (K, Rb, Cs) stabilize larger anions ($O_2^-$) better.
Final Answer: Na → peroxide; K → superoxide.
The most reactive alkali metal with water:
Flame color of K:
The metal that dissolves in liquid NH₃ to give blue solution:
Order of IE in Group 1:
K + O₂ excess gives:
Important Compounds and Anomalous Behavior of LiTopic 2
Important Compounds of Sodium:
1. Sodium Carbonate (Washing Soda, Na₂CO₃·10H₂O):
Solvay Process (industrial preparation):
- $NH_3 + H_2O + CO_2 \to NH_4HCO_3$
- $NH_4HCO_3 + NaCl \to NaHCO_3 + NH_4Cl$
- $2NaHCO_3 \xrightarrow{\Delta} Na_2CO_3 + H_2O + CO_2$
- $2NH_4Cl + Ca(OH)_2 \to 2NH_3 + CaCl_2 + 2H_2O$ (NH₃ recycled)
Uses: Glass manufacture, soaps and detergents, paper, water softening.
2. Sodium Bicarbonate (Baking Soda, NaHCO₃):
- Used in baking (releases CO₂ on heating)
- Mild antiseptic; antacid
- Fire extinguishers
3. Sodium Hydroxide (Caustic Soda, NaOH):
- Made by electrolysis of brine (Castner-Kellner / membrane cell)
- Strong base; used in soap, paper, textile industry
4. Sodium Chloride (NaCl):
- From sea water by evaporation
- Used in food, manufacture of Na, Cl₂, NaOH
Anomalous Properties of Lithium (Diagonal Relationship with Mg):
Li differs from other alkali metals due to:
- Smallest size, highest IE in group
- Highest polarizing power
- Highest hydration energy
Differences from other Group 1 elements:
| Property | Li | Other alkali metals |
|---|---|---|
| Hardness | Hardest | Soft |
| MP | High ($180°$C) | Low |
| Reaction with N₂ | Forms $Li_3N$ (nitride) | Other alkali metals don't form nitride directly |
| Oxide | $Li_2O$ (normal oxide only) | Na → peroxide, K → superoxide |
| LiOH | Decomposes on heating to Li₂O | NaOH, KOH stable |
| Salts | $LiCl$ deliquescent (covalent); $LiF$ insoluble | Other halides ionic, soluble |
| Carbonate | $Li_2CO_3$ unstable; decomposes on heating | Na, K carbonates stable |
Diagonal Relationship Li-Mg:
| Property | Li | Mg |
|---|---|---|
| Reaction with N₂ | $6Li + N_2 \to 2Li_3N$ | $3Mg + N_2 \to Mg_3N_2$ |
| Hydroxide solubility | Less soluble | Less soluble |
| Bicarbonate | LiHCO₃ doesn't exist as solid | Mg(HCO₃)₂ stable in solution only |
| Reaction with water | Slow | Slow |
| Hardness | Hard for alkali metal | Hard |
Explain Solvay process for $Na_2CO_3$ manufacture.
Show solution
Steps:
- $CO_2$ + $NH_3$ + $H_2O$ → $NH_4HCO_3$
- $NH_4HCO_3$ + NaCl → $NaHCO_3$ ↓ + $NH_4Cl$
- $NaHCO_3$ heated → $Na_2CO_3$ + $H_2O$ + $CO_2$
- $NH_4Cl$ + $Ca(OH)_2$ → $NH_3$ + $CaCl_2$ + $H_2O$
$NH_3$ and $CO_2$ are recycled.
Final Answer: Cyclic process using brine, ammonia, lime.
Why is LiCl deliquescent but NaCl is not?
Show solution
Li⁺ has high polarizing power (small, $+1$). LiCl has appreciable covalent character (Fajan's rules). The strong hydration ability of Li⁺ makes LiCl absorb water from air (deliquescent). Na⁺ is bigger; NaCl is purely ionic; not deliquescent.
Final Answer: Small Li⁺ → high hydration energy → deliquescent.
Solvay process produces:
Lithium reacts with N₂ at room temperature to give:
Diagonal relationship of Li is with:
Most ionic alkali halide:
NaHCO₃ on heating gives:
Alkaline Earth Metals (Group 2)
General Trends and PropertiesTopic 1
Alkaline Earth Metals: Group 2 — Be, Mg, Ca, Sr, Ba, Ra. Outer config: $ns^2$. Less reactive than alkali metals; form $M^{2+}$ ions.
Electronic Configuration:
| Element | $Z$ | Config |
|---|---|---|
| Be | 4 | $[He]\,2s^2$ |
| Mg | 12 | $[Ne]\,3s^2$ |
| Ca | 20 | $[Ar]\,4s^2$ |
| Sr | 38 | $[Kr]\,5s^2$ |
| Ba | 56 | $[Xe]\,6s^2$ |
| Ra | 88 | $[Rn]\,7s^2$ (radioactive) |
Trends Down the Group:
| Property | Trend |
|---|---|
| Atomic/Ionic radius | Increases ↑ |
| Sum of first two IEs | Decreases ↓ |
| EN | Decreases ↓ |
| Reactivity with water | Increases (Be doesn't, Mg slow, Ca and below react) |
| MP, BP | Irregular; generally higher than alkali metals (stronger metallic bond from $ns^2$) |
| Hardness | Decreases ↓ |
| Density | Increases ↑ (slowly) |
Properties Compared to Alkali Metals:
- Higher MP, BP
- Harder (stronger metallic bond)
- Higher density
- Higher IE (need to remove 2 electrons)
- Lower reducing power (smaller, harder to lose 2 e⁻)
Flame Colors:
- Be, Mg: don't impart color (small size, high IE, transitions in UV)
- Ca: brick red
- Sr: crimson red
- Ba: apple green
Chemical Reactions:
- With water:
- Be: no reaction
- Mg: with hot water/steam: $Mg + H_2O \to MgO + H_2$
- Ca, Sr, Ba: react with cold water (less vigorously than alkali)
- With air: Burn forming oxides: $2Mg + O_2 \to 2MgO$; Ba forms peroxide $BaO_2$.
- With halogens: $M + X_2 \to MX_2$
- With hydrogen: Form ionic hydrides except $BeH_2$ (covalent, polymeric)
- With acids: $M + 2HCl \to MCl_2 + H_2$
- With nitrogen: Form nitrides $M_3N_2$ (3M + N₂ → $M_3N_2$) at high $T$
- Reducing nature: Less than alkali metals; Ba most reactive in group
Arrange in increasing order of reactivity with water: Be, Mg, Ca, Ba.
Show solution
Be < Mg < Ca < Ba (reactivity increases down the group due to decreasing IE and easier $M^{2+}$ formation).
Final Answer: Be < Mg < Ca < Ba.
Why do Group 2 metals have higher MP than Group 1?
Show solution
Group 2 metals have $ns^2$ configuration, providing 2 valence electrons (per atom) for metallic bonding. Group 1 has only 1. More valence electrons → stronger metallic bond → higher MP, BP, hardness.
Final Answer: $ns^2$ gives stronger metallic bonding than $ns^1$.
Which alkaline earth metal does NOT impart color to flame?
Most reactive alkaline earth metal with water:
Alkaline earth metals form ion:
Be does NOT react with water because:
Density of group 2 elements:
Compounds, Anomalous Behavior of Be, and Diagonal RelationshipTopic 2
Important Compounds:
1. Calcium Oxide (Quicklime, CaO):
- From limestone: $CaCO_3 \xrightarrow{\Delta} CaO + CO_2$
- With water: $CaO + H_2O \to Ca(OH)_2$ (slaked lime)
- Uses: cement, mortar, calcium carbide manufacture
2. Calcium Hydroxide (Slaked Lime, Ca(OH)₂):
- $Ca(OH)_2$ + Cl₂ → bleaching powder $Ca(OCl)Cl$
- Limewater test for CO₂: turns milky due to $CaCO_3$ formation
- Whitewashing
3. Calcium Carbonate (CaCO₃):
- Limestone, chalk, marble
- Manufacture of cement
- Acid neutralizer (antacid)
4. Calcium Sulphate Hemihydrate ($Plaster of Paris, CaSO_4·\frac{1}{2}H_2O$):
- $2CaSO_4·2H_2O \xrightarrow{120°C} 2CaSO_4·\frac{1}{2}H_2O + 3H_2O$
- Sets with water to form gypsum
5. Cement: Calcium silicate + aluminate; produced from limestone + clay + gypsum.
6. Magnesium Hydroxide:
- Milk of magnesia — mild antacid
Anomalous Behavior of Beryllium (Diagonal Relationship with Al):
Be differs from other alkaline earth metals because:
- Smallest size
- Highest IE
- Higher EN
Properties of Be unique to Group 2:
| Property | Be | Other Group 2 |
|---|---|---|
| Salts | Covalent (e.g., BeCl₂, BeSO₄) | Ionic |
| Reactivity with water | No reaction | Some react with water |
| Oxide | Amphoteric ($BeO$ reacts with both acid and base) | Basic |
| Carbide | $Be_2C$ gives methane on hydrolysis | $CaC_2$ gives acetylene |
| Hydroxide | $Be(OH)_2$ amphoteric | Basic |
| Coordination number | $4$ (small size) | $6$ |
| Halides | Dimeric $Be_2Cl_4$, polymeric | Monomeric |
Diagonal Relationship Be-Al:
| Property | Be | Al |
|---|---|---|
| Reaction with NaOH | Yes: $Be + 2NaOH \to Na_2BeO_2 + H_2$ | Yes: $2Al + 2NaOH + 2H_2O \to 2NaAlO_2 + 3H_2$ |
| Carbides | $Be_2C \to CH_4$ on hydrolysis | $Al_4C_3 \to CH_4$ |
| Halides | $BeCl_2$ dimeric/polymeric (3c-4e bonds) | $AlCl_3$ dimeric $Al_2Cl_6$ |
| Hydroxide | Amphoteric | Amphoteric |
| Atom small, charge density high | Yes | Yes |
Why is $BeO$ amphoteric while $MgO$ is basic?
Show solution
Be has very high polarizing power (small, $+2$). $BeO$ has significant covalent character; the O-Be bond is more like polar covalent. So it can react with both acids ($BeO + 2HCl \to BeCl_2 + H_2O$) and bases ($BeO + 2NaOH \to Na_2BeO_2 + H_2O$). $MgO$ is more ionic; reacts only with acids — purely basic.
Final Answer: Be high polarizing power → covalent BeO → amphoteric.
Predict products: $CaC_2 + H_2O \to ?$ vs $Be_2C + H_2O \to ?$
Show solution
$CaC_2 + 2H_2O \to Ca(OH)_2 + C_2H_2$ (acetylene) $Be_2C + 4H_2O \to 2Be(OH)_2 + CH_4$ (methane)
The carbide ion in CaC₂ is $C_2^{2-}$ (acetylide); in Be₂C it is $C^{4-}$ (methide).
Final Answer: CaC₂ → acetylene; Be₂C → methane.
Plaster of Paris is:
Quicklime is:
Diagonal relationship of Be is with:
Beryllium hydroxide is:
Limewater turns milky with:
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