1. General Characteristics

The s-block elements are those in which the last electron enters the outermost s-orbital. They are situated on the extreme left of the periodic table.

General Electronic Configuration:

Group 1 (Alkali Metals): [Noble Gas] ns¹
Group 2 (Alkaline Earth Metals): [Noble Gas] ns²

2. Alkali Metals (Group 1)

Elements: Li, Na, K, Rb, Cs, Fr. Called "Alkali" because they form strong hydroxides with water.

Physical Properties

Atomic Size: Largest in their respective periods. Increases down the group.
Hydration Enthalpy: Decreases down the group.
Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺
(Li⁺ has max degree of hydration, hence Li salts are hydrated e.g., LiCl·2H₂O).
Flame Color: Valence electrons get excited and emit color.
Li (Crimson), Na (Yellow), K (Violet), Rb (Red violet), Cs (Blue).

Chemical Properties

1. Reactivity with Air (Oxides):

They burn vigorously to form oxides.
• Lithium forms Monoxide (Li₂O).
• Sodium forms Peroxide (Na₂O₂).
• K, Rb, Cs form Superoxides (MO₂).

2. Reactivity with Liquid Ammonia:

Dissolve to give deep blue solution (conducting).
M + (x+y)NH₃ → [M(NH₃)_x]⁺ + [e(NH₃)_y]^-
Blue color is due to Ammoniated Electron. Paramagnetic.

Important Compounds

Sodium Carbonate (Washing Soda): Na₂CO₃·10H₂O. Prepared by Solvay Process.
Sodium Hydroxide (Caustic Soda): NaOH. Prepared by electrolysis of brine (Castner-Kellner cell).
Sodium Bicarbonate (Baking Soda): NaHCO₃. Decomposes on heating to give CO₂ (makes cakes fluffy).

3. Alkaline Earth Metals (Group 2)

Elements: Be, Mg, Ca, Sr, Ba, Ra. Their oxides/hydroxides are alkaline and found in earth's crust.

Physical Properties

Ionization Enthalpy: Higher than Group 1 (due to smaller size). Decreases down the group.
Flame Color:
Ca (Brick Red), Sr (Crimson), Ba (Apple Green).
Be and Mg do NOT impart color (electrons held too tightly).

Solubility Trends

Compound	Trend Down the Group	Reason
Hydroxides	Solubility Increases	Lattice Enthalpy decreases more than Hydration Enthalpy.
Sulphates	Solubility Decreases	Hydration Enthalpy decreases rapidly.

Important Compounds of Calcium

Quick Lime (CaO): From heating limestone. Basic oxide.
Slaked Lime (Ca(OH)₂): CaO + H₂O. Used in whitewashing. Aq. solution is Lime Water.
Plaster of Paris (POP): CaSO₄·½H₂O. Obtained by heating Gypsum (CaSO₄·2H₂O) at 373 K.

4. Anomalous & Diagonal Relationships

Anomalous Behavior of Lithium:

Due to exceptionally small size and high polarizing power.
• Li is harder than other alkali metals.
• Li forms nitride (Li₃N) directly with N₂.
• LiNO₃ decomposes to oxide (Li₂O), others give nitrite (MNO₂).

Diagonal Relationship (Li - Mg):

Li and Mg show similar properties due to similar charge/size ratio.
• Both form nitrides.
• Carbonates decompose on heating.
• Chlorides (LiCl, MgCl₂) are deliquescent and soluble in ethanol.

Structure of BeCl₂

Solid State: Polymeric chain structure with chloro-bridges.
Vapor Phase: Dimer (Be₂Cl₄) below 1200 K, Linear monomer (BeCl₂) above 1200 K.

5. Biological Importance

Sodium (Na⁺) & Potassium (K⁺):

Na⁺: Major cation in Extracellular fluid. Regulates blood pressure, nerve signals.
K⁺: Major cation in Intracellular fluid. Activates enzymes, nerve signals.
Na-K Pump: Consumes 1/3rd of ATP in resting animal.

Magnesium (Mg²⁺) & Calcium (Ca²⁺):

Mg²⁺: Cofactor for enzymes (ATP). Central atom in Chlorophyll (photosynthesis).
Ca²⁺: 99% in bones/teeth (Apatite). Required for muscle contraction and blood clotting.

Numericals & HOTS

Q1. Oxidation States

Calculate the oxidation number of Oxygen in Potassium Superoxide (KO₂) and explain why the molecule is paramagnetic.

Solution:
1. Potassium (Group 1) always has O.N. = +1.
Let O.N. of Oxygen be x.
(+1) + 2x = 0 → 2x = -1 → x = -1/2.

2. Paramagnetism: The superoxide ion (O₂^-) has an odd number of electrons (17e^-). According to Molecular Orbital Theory, it has one unpaired electron in the π* antibonding orbital, making it paramagnetic.

Q2. Solubility of Ca(OH)₂

The pH of a saturated solution of Ca(OH)₂ is 9. Calculate the solubility product (K_sp) of Ca(OH)₂.

Solution:
pH = 9 → pOH = 14 - 9 = 5.
[OH^-] = 10^-5 M.

Dissociation: Ca(OH)₂ ↔ Ca²⁺ + 2OH^-
If solubility is 's', then [Ca²⁺] = s and [OH^-] = 2s.
Given 2s = 10^-5 → s = 0.5 × 10^-5.

K_sp = [Ca²⁺][OH^-]²
K_sp = (0.5 × 10^-5) × (10^-5)²
K_sp = 0.5 × 10^-15
K_sp = 5 × 10^-16

Q3. Lithium Nitrate Decomposition

Calculate the total volume of gases evolved at STP when 13.8 g of LiNO₃ is heated. (Li=7, N=14, O=16).

Solution:
Reaction: 4LiNO₃ → 2Li₂O + 4NO₂ + O₂.
Molar Mass LiNO₃ = 7 + 14 + 48 = 69 g/mol.
Moles taken = 13.8 / 69 = 0.2 mol.

From stoichiometry:
4 mol LiNO₃ gives (4 + 1) = 5 mol gas.
0.2 mol LiNO₃ gives (5/4) × 0.2 = 0.25 mol gas.

Volume = n × 22.4 L = 0.25 × 22.4
Volume = 5.6 Litres

Q4. Hydration vs Mobility (HOTS)

Why is the ionic mobility of Cs⁺ higher than Li⁺ in aqueous solution, despite Li⁺ being smaller?

Solution:
Concept: Hydration Enthalpy.
Small ions have high charge density and attract more water molecules (High Hydration).

Li⁺ (smallest) becomes heavily hydrated → [Li(H₂O)_n]⁺.
The Effective Size of hydrated Li⁺ becomes very large, making it move slowly.
Cs⁺ (largest) is least hydrated, so its effective size remains small, allowing it to move faster.
Mobility: Cs⁺ > Rb⁺ > K⁺ > Na⁺ > Li⁺

Q5. Percentage Purity

20 g of a limestone sample (CaCO₃) on heating gives 8.4 g of CaO. Calculate the percentage purity of CaCO₃.

Solution:
Reaction: CaCO₃ → CaO + CO₂.
Molar Mass: CaCO₃ = 100g, CaO = 56g.

100 g CaCO₃ produces 56 g CaO.
To produce 8.4 g CaO, pure CaCO₃ needed = (100/56) × 8.4 = 15 g.

% Purity = (Pure Mass / Sample Mass) × 100
% Purity = (15 / 20) × 100
Ans: 75%

Q6. Lattice vs Hydration (HOTS)

Explain why MgSO₄ is soluble in water while BaSO₄ is practically insoluble.

Solution:
Solubility depends on the balance between Lattice Enthalpy and Hydration Enthalpy.

For Sulphates (large anion SO₄^2-), Lattice Enthalpy remains almost constant down the group.
However, Hydration Enthalpy decreases significantly from Mg²⁺ to Ba²⁺.

For MgSO₄: Hydration Energy > Lattice Energy (Soluble).
For BaSO₄: Hydration Energy < Lattice Energy (Insoluble).

Q7. Gypsum Conversion

Calculate the mass of water lost when 172 g of Gypsum is heated to 393 K to form Plaster of Paris. (MW Gypsum = 172).

Solution:
Reaction: CaSO₄·2H₂O → CaSO₄·½H₂O + 1.5H₂O.
Moles of Gypsum = 172 / 172 = 1 mol.

Water lost per mole = 1.5 moles.
Mass of water = 1.5 × 18 g/mol
Ans: 27 g

Q8. Recycling in Solvay (HOTS)

In the Solvay process, how is Ammonia recovered, and what is the only waste product generated?

Solution:
Recovery: The filtrate containing NH₄Cl is treated with Ca(OH)₂ (Slaked Lime).
2NH₄Cl + Ca(OH)₂ → 2NH₃ + CaCl₂ + 2H₂O.
Ammonia is reused.

Waste Product: Calcium Chloride (CaCl₂).

Q9. Identifying Compounds

Alkali metal 'A' reacts with Nitrogen to form 'B'. 'A' reacts with O₂ to form 'C' which is a normal oxide. Identify A, B, and C.

Solution:
Only Lithium reacts directly with Nitrogen.
So, A = Li.

Reaction with N₂: 6Li + N₂ → 2Li₃N.
So, B = Li₃N (Lithium Nitride).

Reaction with O₂: 4Li + O₂ → 2Li₂O.
So, C = Li₂O (Lithium Oxide).

Q10. Amphoteric Reactions (HOTS)

Write reactions to show the amphoteric nature of Beryllium Hydroxide.

Solution:
Amphoteric means reacting with both acid and base.

With Acid (HCl):
Be(OH)₂ + 2HCl → BeCl₂ + 2H₂O (Acts as Base).

With Base (NaOH):
Be(OH)₂ + 2NaOH → Na₂[Be(OH)₄] (Acts as Acid).
Product: Sodium Beryllate.

Important Trends & Formulae

1. General Physical Trends

Hydration Enthalpy (Size of Hydrated Ion):

Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺

Be²⁺ > Mg²⁺ > Ca²⁺ > Sr²⁺ > Ba²⁺

Ionic Mobility (in water):

Cs⁺ > Rb⁺ > K⁺ > Na⁺ > Li⁺

(Smaller hydrated ions move faster)

2. Solubility Trends (Group 2)

Hydroxides (Increases down):

Be(OH)₂ < Mg(OH)₂ < Ca(OH)₂ < Sr(OH)₂ < Ba(OH)₂

Sulphates (Decreases down):

BeSO₄ > MgSO₄ > CaSO₄ > SrSO₄ > BaSO₄

(Based on Lattice Energy vs Hydration Energy)

3. Thermal Stability

Carbonates & Nitrates:

Increases down the group

Decomposition of LiNO₃:

4LiNO₃ → 2Li₂O + 4NO₂ + O₂

(Other alkali nitrates give Nitrite: 2NaNO₃ → 2NaNO₂ + O₂)

4. Common Chemical Formulae

Washing Soda	Na₂CO₃ · 10H₂O
Baking Soda	NaHCO₃
Glauber's Salt	Na₂SO₄ · 10H₂O
Epsom Salt	MgSO₄ · 7H₂O
Gypsum	CaSO₄ · 2H₂O
POP	CaSO₄ · ½H₂O

20 Golden Facts (NEET)

1. Density Anomaly: Generally, density increases down the group. However, Potassium (K) is lighter than Sodium (Na) due to an unusual increase in atomic size of K.
2. Lithium's Nitride: Lithium is the only alkali metal that reacts directly with Nitrogen gas to form a Nitride (Li₃N). This is due to its small size and high lattice energy.
3. Superoxides: Potassium (K), Rubidium (Rb), and Cesium (Cs) form superoxides (MO₂). KO₂ is used in submarines and space shuttles to absorb CO₂ and release O₂.
4. Solution in Ammonia: All alkali metals dissolve in liquid ammonia giving a Deep Blue solution which is conducting and paramagnetic (due to ammoniated electrons). On standing, it liberates H₂ and becomes amide (Blue → Bronze).
5. Photoelectric Effect: Cesium (Cs) and Potassium (K) have very low ionization enthalpies, making them useful as electrodes in photoelectric cells.
6. Solvay Process Limit: Solvay process is used to manufacture Na₂CO₃. It cannot represent manufacture of K₂CO₃ because KHCO₃ is too soluble to be precipitated.
7. Flame Colors:
Li: Crimson Red | Na: Golden Yellow
K: Lilac (Violet) | Rb: Red Violet | Cs: Blue
Ca: Brick Red | Sr: Crimson | Ba: Apple Green
8. No Flame Color: Beryllium (Be) and Magnesium (Mg) atoms are small and their valence electrons are strongly bound. The energy of the flame is insufficient to excite them; hence, they show no flame color.
9. Sorel's Cement: A mixture of MgO and MgCl₂ solution forms a hard setting cement known as Sorel's cement (Magnesium oxychloride).
10. Amphoteric Be: Beryllium hydroxide [Be(OH)₂] and Oxide (BeO) are Amphoteric in nature (react with both acid and alkali). Other Group 2 oxides are basic.
11. Dead Burnt Plaster: If Gypsum is heated above 393 K, all water of crystallization is lost, forming anhydrous CaSO₄, known as "dead burnt plaster" (it does not set with water).
12. Castner-Kellner Cell: Used for NaOH production. Mercury cathode is used, forming Na-Hg amalgam, which prevents the reaction of Na with water in the cell.
13. Solubility of Li Salts: LiF is least soluble (high lattice energy). LiI is soluble in ethanol/acetone (high covalent character due to polarization).
14. BeCl₂ Structure: In solid state, it has a polymeric chain structure with chloro-bridges. In vapor phase, it exists as a dimer (Be₂Cl₄) or monomer (linear BeCl₂ at high temp).
15. Reducing Power: In aqueous solution, Lithium is the strongest reducing agent (most negative E° = -3.04 V) due to its high hydration energy.
16. Biological Role: A typical 70 kg man contains ~170 g Ca (bones) but only ~5 g Fe. Mg is present in Chlorophyll (light absorption).
17. Suspension of Lime: An aqueous suspension of slaked lime is called Milk of Lime. The clear solution is called Lime Water.
18. Barium Meal: BaSO₄ is insoluble and opaque to X-rays. It is used as a "Barium meal" for X-ray imaging of the digestive tract.
19. Thermal Stability Order: BeCO₃ < MgCO₃ < CaCO₃ < SrCO₃ < BaCO₃. BeCO₃ is unstable and kept in an atmosphere of CO₂.
20. Diagonal Relationship: Lithium shows diagonal relationship with Magnesium. Beryllium shows diagonal relationship with Aluminium. (Similar charge/size ratio).

📱 Practice MCQs for this topic inside our App

1. General Characteristics

2. Alkali Metals (Group 1)

Physical Properties

Chemical Properties

Important Compounds

3. Alkaline Earth Metals (Group 2)

Physical Properties

Solubility Trends

Important Compounds of Calcium

4. Anomalous & Diagonal Relationships

Structure of BeCl2

5. Biological Importance

Numericals & HOTS

Important Trends & Formulae

20 Golden Facts (NEET)

Structure of BeCl₂