Alkaline Earth Metals (Group 2) (Beryllium, Magnesium, Calcium, Strontium, Barium, Radium) : Properties, Trends and Reactions Explained For CSIR-NET, IIT-JAM, BITSAT, GATE, JEE Advanced Exams

Group 2 — The Alkaline Earth Elements: A Complete Study Guide

JEE Advanced NEET IIT-JAM BITSAT GATE CSIR-NET

Imagine the periodic table as a grand apartment building. Group 1 elements (alkali metals) are the wild, reactive tenants on the left — they explode in water, they melt at almost room temperature, and they are outrageously reactive. Right next door in Group 2 live their cousins — the alkaline earth metals: Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba), and Radium (Ra). These neighbours are reactive too, but calmer, denser, harder, and far more useful industrially. This chapter is a goldmine for competitive exams, and we are going to explore every corner of it — from electronic configuration to Grignard reagents — in crystal-clear language with diagrams, reaction equations, and exam-ready tricks.

📌 Why This Chapter Matters in Competitive Exams Group 2 chemistry is one of the highest-yield chapters across JEE, NEET, GATE, and CSIR-NET. Expect 2–5 direct questions every year from trends in properties, thermal stability of carbonates/sulphates, anomalous behaviour of Be, hard water chemistry, and Grignard reagents.

1. Electronic Configuration — The Root of Everything

Every Group 2 element has two electrons in its outermost s-orbital. This single fact explains almost all of their chemical behaviour — why they form M²⁺ ions, why they are harder than Group 1 metals, and why they burn so brilliantly in air.

Element	Symbol	Electronic Configuration	Short Form
Beryllium	Be	1s² 2s²	[He] 2s²
Magnesium	Mg	1s² 2s² 2p⁶ 3s²	[Ne] 3s²
Calcium	Ca	1s² 2s² 2p⁶ 3s² 3p⁶ 4s²	[Ar] 4s²
Strontium	Sr	...... 4s² 4p⁶ 5s²	[Kr] 5s²
Barium	Ba	...... 5s² 5p⁶ 6s²	[Xe] 6s²
Radium	Ra	...... 6s² 6p⁶ 7s²	[Rn] 7s²

🎯 Exam Tip — JEE/NEET Calcium's configuration ends at 4s², not 3d. Students often confuse Ca with transition metals. Ca has Z=20; electrons fill 1s→2s→2p→3s→3p→4s. The 3d orbitals remain empty. This makes Ca a representative element, not a transition metal.

2. Occurrence and Abundance in Nature

These metals are called "earth" metals partly because many of their oxides were historically called "earths" — and these oxides are quite stable (they don't melt easily). Let us see where they hide in nature:

Beryllium (Be): Rare — only about 2 ppm in the Earth's crust. Found in the silicate mineral beryl (Be₃Al₂Si₆O₁₈) and phenacite (Be₂SiO₄). The gemstone emerald is beryl with traces of chromium giving it the green colour.
Magnesium (Mg): 6th most abundant element in the crust (~27,640 ppm). Occurs as dolomite [MgCO₃·CaCO₃], magnesite (MgCO₃), epsomite (MgSO₄·7H₂O). Present in seawater as Mg²⁺ ions (~0.13%).
Calcium (Ca): 5th most abundant element (~46,600 ppm). Massive deposits of CaCO₃ as limestone, marble, chalk, and coral. Also found as gypsum (CaSO₄·2H₂O) and fluorite (CaF₂).
Strontium (Sr): 384 ppm. Found as celestite (SrSO₄) and strontianite (SrCO₃).
Barium (Ba): 390 ppm. Found primarily as barytes (BaSO₄).
Radium (Ra): Extremely scarce (~1.3 × 10^-6 ppm) and radioactive. Isolated from uranium ore (pitchblende) by Pierre and Marie Curie. Marie Curie received the Nobel Prize in Chemistry in 1911 for this work.

🧠 Mnemonic — Order of Abundance Mg and Ca are far more abundant than Sr, Ba, Ra. Think: "More Mg, Colossal Ca, Sparse Sr, Barely Ba, Rare Ra."

3. Physical Properties — Size, Density, Melting Points

3.1 Atomic and Ionic Size

As we go down Group 2, more electron shells are added, so both atomic radius and ionic radius increase. However, all Group 2 atoms are smaller than their Group 1 neighbours in the same period. Why? Because the nucleus has one extra proton, pulling the electrons inward more tightly.

Element	Metallic Radius (Å)	Ionic Radius M²⁺ (Å)	Density (g cm^-3)
Be	1.12	0.31	1.85
Mg	1.60	0.72	1.74
Ca	1.97	1.00	1.55
Sr	2.15	1.18	2.63
Ba	2.22	1.35	3.62
Ra	—	1.48	~5.5

🎯 Key Observation for GATE/JAM The jump in ionic size from Be²⁺ (0.31 Å) to Mg²⁺ (0.72 Å) is four times greater than the jump from Li⁺ to Na⁺. This enormous size difference is why beryllium behaves so differently from the rest of the group. Note: Ca shows a slight density dip compared to Mg — this is because Ca adopts a different crystal structure (face-centred cubic).

Figure 1: Ionic radius (M²⁺) increases steadily down Group 2 as more electron shells are added.

3.2 Melting and Boiling Points

Group 2 metals have two valence electrons participating in metallic bonding, compared to one for Group 1. This gives them higher cohesive energy — they are harder, denser, and have much higher melting and boiling points. However, the melting points do not follow a perfectly regular trend because different metals adopt different crystal structures.

Element	Melting Point (°C)	Boiling Point (°C)
Be	1287	~2500
Mg	649	1105
Ca	839	1494
Sr	768	1381
Ba	727	~1850

4. Ionisation Energy and Electronegativity

The third ionisation energy for all Group 2 elements is extremely high — so high that M³⁺ ions are never formed under ordinary conditions. The elements always lose exactly two electrons to give M²⁺.

Element	IE₁ (kJ mol^-1)	IE₂ (kJ mol^-1)	IE₃ (kJ mol^-1)	Electronegativity
Be	899	1757	14847	1.5
Mg	737	1450	7731	1.2
Ca	590	1145	4910	1.0
Sr	549	1064	—	1.0
Ba	503	965	—	0.9

The total energy to produce M²⁺ (IE₁ + IE₂) is four times greater than the energy needed to produce M⁺ from Group 1. Despite this, ionic compounds are formed because the energy released when the crystal lattice forms more than compensates.

🎯 JEE Advanced / GATE Classic Beryllium's electronegativity (1.5) is significantly higher than the others. This means the electronegativity difference between Be and other elements (like F, O, Cl) is often too small to form truly ionic bonds. Hence, Be compounds are predominantly covalent — this is the key to understanding anomalous behaviour of Be.

5. Hydration Energies — Why Group 2 Salts Hold More Water

Group 2 ions are smaller and carry a 2+ charge. This means they attract water molecules with immense force — their hydration energies are four to five times greater than Group 1 ions.

Ion	Ionic Radius (Å)	ΔH_hydration (kJ mol^-1)
Be²⁺	0.31	−2494
Mg²⁺	0.72	−1921
Ca²⁺	1.00	−1577
Sr²⁺	1.18	−1443
Ba²⁺	1.35	−1305

This is why Group 2 chlorides have water of crystallisation: MgCl₂·6H₂O, CaCl₂·6H₂O, BaCl₂·2H₂O — while Group 1 chlorides NaCl and KCl are anhydrous. The number of water molecules decreases as ions get larger, because larger ions hold water less tightly.

6. The Anomalous Behaviour of Beryllium — The Oddball of Group 2

Beryllium is the rebellious younger sibling of Group 2. While Mg, Ca, Sr, and Ba behave predictably as ionic metals, Be marches to its own tune — forming covalent compounds, polymeric structures, and showing a diagonal relationship with aluminium. Here is exactly why:

Three Reasons for Be's Uniqueness

Extremely small size: The Be²⁺ ion has a radius of only 0.31 Å — the smallest among all Group 2 ions. According to Fajans' rules, a small, highly charged cation can distort the electron cloud of the anion so effectively that the bond gains covalent character. Be²⁺ does exactly this — it polarises anions strongly and forms covalent bonds.
High electronegativity (1.5): When Be reacts with another element, the electronegativity difference is often not large enough for a fully ionic bond. Even BeF₂ (where the difference is 2.5) shows some covalent character; BeO (difference 2.0) is clearly covalent in character.
Only four orbitals available (second period): Be is in the second row of the periodic table. Its outer shell can hold only eight electrons maximum, using one 2s and three 2p orbitals. The maximum conventional electron-pair bond count is four, giving Be a maximum coordination number of 4 in most compounds. Larger Group 2 elements like Mg can use 3d orbitals as well and achieve coordination number 6.

Figure 2: BeCl₂ exists as (a) linear monomer in gas phase, (b) dimer with chlorine bridges, and (c) chain polymer in the solid state. Bridging Cl atoms use a lone pair to form a coordinate bond to a second Be atom.

Diagonal Relationship: Be and Al

Beryllium in Group 2 behaves remarkably like aluminium in Group 13 — they are diagonal neighbours in the periodic table. The reason is that their atomic and ionic sizes are similar. Here is a side-by-side comparison:

Property	Be (Group 2)	Al (Group 13)
Compounds character	Predominantly covalent	Predominantly covalent
Halides	Polymeric, electron-deficient	Dimeric (Al₂Cl₆), electron-deficient
Hydroxide	Be(OH)₂ — amphoteric	Al(OH)₃ — amphoteric
Reaction with NaOH	Forms beryllate + H₂	Forms aluminate + H₂
Passive with conc. HNO₃?	Yes	Yes
Carbide type	Be₂C (methanide, gives CH₄)	Al₄C₃ (methanide, gives CH₄)
Standard electrode potential	−1.85 V	−1.66 V (close!)

7. Chemical Properties — Reactivity Trends

7.1 Reaction with Water

The reactivity of Group 2 metals with water increases down the group, but the pattern is irregular at the top:

Be: Does not react with water even at high temperatures. Possibly reacts with steam to form BeO, but this is uncertain. Be's low reduction potential (−1.85 V) makes it much less electropositive than the others.
Mg: Does not react with cold water (a protective oxide layer forms, resembling aluminium). Slowly reacts with hot water or steam to give Mg(OH)₂ or MgO.
Ca, Sr, Ba: React readily with cold water, liberating hydrogen gas and forming metal hydroxides.

Ca + 2H₂O → Ca(OH)₂ + H₂↑

Mg + H₂O (hot) → MgO + H₂↑

Ba + 2H₂O → Ba(OH)₂ + H₂↑

7.2 Reaction with Acids

All Group 2 metals react with dilute acids (HCl, H₂SO₄) to liberate hydrogen gas. Beryllium, however, is rendered passive by concentrated HNO₃ — a thin oxide layer forms that protects the surface from further attack (same as aluminium).

M + 2HCl → MCl₂ + H₂↑

Be + 2NaOH + 2H₂O → Na₂[Be(OH)₄] + H₂↑

(Be is amphoteric — reacts with both acids and alkalis)

7.3 Reaction with Oxygen

All Group 2 metals burn in oxygen to form normal oxides (MO). Barium also forms a peroxide (BaO₂) when excess oxygen is present. Magnesium burns in air with such dazzling brilliance that it is used to start thermite reactions and historically in flash photography.

2Mg + O₂ → 2MgO (dazzling white flame)

2Ba + O₂ → 2BaO (normal oxide)

Ba + O₂ (excess) → BaO₂ (barium peroxide)

Mg + air → MgO + Mg₃N₂ (mixture)

🎯 JEE/NEET Trick — Which elements form peroxides? In Group 2, only Ba forms a peroxide (BaO₂) under ordinary conditions. SrO₂ needs high pressure. CaO₂ can be made as a hydrate but not directly. No peroxide of Be is known. This contrasts with Group 1 where Na forms Na₂O₂, K forms KO₂ (superoxide), Li forms only Li₂O.

8. Hydroxides — Basic Strength and Solubility Trend

The hydroxides show a perfectly regular trend: basic strength increases from Be to Ba. This is because as the metal ion gets larger, it holds the OH⁻ less tightly — the M–O bond weakens, making it easier for OH⁻ to dissociate in water.

Be(OH)₂: Amphoteric — dissolves in both acid and alkali. Weakly acidic character due to high charge density of Be²⁺.
Mg(OH)₂: Weakly basic — sparingly soluble (~1 × 10^-4 g/L). Used as an antacid ("milk of magnesia").
Ca(OH)₂: Moderately strong base, solubility ~2 g/L. Called slaked lime. Its solution is called lime water.
Sr(OH)₂: Strong base, solubility ~8 g/L.
Ba(OH)₂: Strong base, solubility ~39 g/L. Its solution is called baryta water.

Be(OH)₂ + 2HCl → BeCl₂ + 2H₂O (acidic behaviour)

Be(OH)₂ + 2NaOH → Na₂[Be(OH)₄] (basic behaviour, amphoteric)

Figure 3: Solubility of Group 2 hydroxides in g/L at 20°C. Note the enormous jump from Ca to Ba — the trend is opposite to that of sulphates.

9. Solubility Patterns — The Opposite Trends of Hydroxides and Sulphates

This is one of the most frequently tested topics in JEE and NEET. The solubility of Group 2 compounds depends on the balance between lattice energy (energy needed to break the crystal) and hydration energy (energy released when ions are hydrated). These two factors change in opposite directions as you go down the group:

Both lattice energy and hydration energy decrease going down (ions get larger).
For most compounds (chlorides, nitrates, sulphates), the hydration energy decreases faster than the lattice energy → solubility decreases down the group.
For fluorides and hydroxides, the lattice energy decreases faster than the hydration energy → solubility increases down the group.

Compound Type	Trend (Be → Ba)	Example
Hydroxides M(OH)₂	Solubility increases ↑	Mg(OH)₂ ≪ Ba(OH)₂
Fluorides MF₂	Solubility increases ↑	BeF₂ (very soluble) → BaF₂ (insoluble)
Sulphates MSO₄	Solubility decreases ↓	BeSO₄ (very soluble) → BaSO₄ (insoluble)
Carbonates MCO₃	Solubility decreases ↓	All sparingly soluble; BaCO₃ least soluble
Chlorides MCl₂	Solubility decreases ↓	All soluble, slightly less down group

🧠 Golden Rule for Exams "Sulphates and Carbonates — Solubility Decreases. Hydroxides and Fluorides — Solubility Increases." Remember: BaSO₄ is so insoluble it is used in X-ray barium meals; Mg(OH)₂ is insoluble but Ba(OH)₂ is quite soluble.

10. Thermal Stability of Carbonates and Sulphates

This is arguably the most tested concept from this chapter in JEE Advanced, GATE, and CSIR-NET. The key principle: the more basic the metal, the more stable its carbonate and sulphate toward heat.

As you go down Group 2, the metal ions get larger. A large cation can stabilise the large CO₃²⁻ and SO₄²⁻ ions through favourable lattice energy. Smaller cations (like Be²⁺, Mg²⁺) prefer to stabilise smaller oxide O²⁻ ions instead — so they decompose carbonates more easily.

Decomposition Temperatures of Carbonates:

BeCO₃ → BeO + CO₂ (<100°C)

MgCO₃ → MgO + CO₂ (540°C)

CaCO₃ → CaO + CO₂ (900°C) ← Industrial quicklime production

SrCO₃ → SrO + CO₂ (1290°C)

BaCO₃ → BaO + CO₂ (1360°C)

Decomposition Temperatures of Sulphates:

BeSO₄ → BeO + SO₃ (500°C)

MgSO₄ → MgO + SO₃ (895°C)

CaSO₄ → CaO + SO₃ (1149°C)

SrSO₄ → SrO + SO₃ (1374°C)

Figure 4: Decomposition temperature of Group 2 carbonates increases sharply from Be to Ba, illustrating that larger metal ions stabilise the carbonate lattice more effectively.

11. Hardness of Water — Temporary and Permanent

Hard water contains dissolved Ca²⁺ and Mg²⁺ ions. These ions react with stearate ions in soap to form an insoluble scum — so no lather forms. Hard water also deposits scale in pipes, boilers, and kettles.

Temporary Hardness

Caused by dissolved Ca(HCO₃)₂ and Mg(HCO₃)₂. Called "temporary" because it can be removed by boiling:

Ca(HCO₃)₂ → CaCO₃↓ + H₂O + CO₂↑

2HCO₃⁻ ⇌ CO₃²⁻ + H₂O + CO₂

Temporary hardness can also be removed by adding slaked lime (lime softening) at pH 10.5:

Ca(HCO₃)₂ + Ca(OH)₂ ⇌ 2CaCO₃↓ + 2H₂O

Permanent Hardness

Caused by CaSO₄ and MgSO₄ in solution. Cannot be removed by boiling. Removal methods include:

Ion-exchange resins (Ca²⁺ and Mg²⁺ replaced by Na⁺)
Adding Na₂CO₃ (Solvay process): CaSO₄ + Na₂CO₃ → CaCO₃↓ + Na₂SO₄
Adding polyphosphates (EDTA, Na₃PO₄) to sequester Ca²⁺ and Mg²⁺
Distillation (laboratory method)

🎯 Cave Chemistry — JEE Favourite Stalactites (from ceiling) and stalagmites (from floor) form because water percolating through limestone contains Ca²⁺(HCO₃⁻)₂. This solution slowly decomposes: Ca²⁺(HCO₃⁻)₂ → CaCO₃↓ + CO₂↑ + H₂O The insoluble CaCO₃ deposits slowly over thousands of years, building these beautiful structures.

12. Detection of CO₂ — Lime Water and Baryta Water Tests

Two important analytical reactions involving Group 2 hydroxides are used to detect carbon dioxide:

Lime water test (Ca(OH)₂ solution):

Ca(OH)₂ + CO₂ → CaCO₃↓ (milky) + H₂O

CaCO₃ + CO₂ + H₂O → Ca(HCO₃)₂ (excess CO₂ clears the solution)

Baryta water test (Ba(OH)₂ solution — more sensitive):

Ba(OH)₂ + CO₂ → BaCO₃↓ (milky) + H₂O

13. Important Oxides and Their Industrial Uses

Calcium Oxide (Quicklime, CaO)

Produced on an enormous scale (127.9 million tonnes in 1993) by roasting limestone in lime kilns:

CaCO₃ → CaO + CO₂ (at ~900°C)

Uses of lime (CaO):

Steel making: removes phosphates and silicates as slag
Cement production: mixed with SiO₂, Al₂O₃, and clay
Glass making
Water softening
Making Ca(OH)₂ (slaked lime) by adding water
Bleaching powder: 3Ca(OH)₂ + 2Cl₂ → Ca(OCl)₂·Ca(OH)₂·CaCl₂·2H₂O

BeO and MgO — Refractory Materials

Both BeO (m.p. ~2500°C) and MgO (m.p. ~2800°C) are excellent refractories (materials that withstand extremely high temperatures in furnaces). Their useful properties include high melting points, very low vapour pressure, excellent thermal conductivity, chemical inertness, and electrical insulation. BeO is also used in nuclear reactors because it has a very low neutron-capture cross-section.

14. Sulphates — From Plaster of Paris to Barium Meals

Group 2 sulphates are among the most practically important compounds in this chapter:

BeSO₄ and MgSO₄: Very soluble (high hydration energy of small ions). Epsom salt is MgSO₄·7H₂O — used as a mild laxative.
CaSO₄·2H₂O (Gypsum): World production was 88.2 million tonnes in 1992. Used in plasterboard, plaster, cement.
Plaster of Paris (CaSO₄·½H₂O): Made by heating gypsum to 150°C. When mixed with water, it sets hard by reabsorbing water to reform gypsum. Used in construction, dental moulds, orthopaedic casts.
BaSO₄: Completely insoluble in water, non-toxic. Used as a "barium meal" — patient drinks a suspension, and since BaSO₄ is opaque to X-rays, it outlines the stomach and duodenum, helping diagnose ulcers.

CaSO₄·2H₂O →^150°C CaSO₄·½H₂O →^200°C CaSO₄ →^1100°C CaO + SO₃

(gypsum) → (plaster of Paris) → (anhydrite) → decomposition

15. Nitrides, Carbides, Hydrides — Completing the Picture

Nitrides

All Group 2 metals burn in nitrogen at high temperatures to form ionic nitrides M₃N₂. This is different from Group 1 where only lithium forms a nitride (Li₃N) under ordinary conditions.

3Ca + N₂ → Ca₃N₂

Ca₃N₂ + 6H₂O → 3Ca(OH)₂ + 2NH₃

Carbides

Calcium carbide (CaC₂) is the most famous. It reacts with water to produce ethyne (acetylene), which was historically used in oxy-acetylene welding:

CaO + 3C →^~2000°C CaC₂ + CO

CaC₂ + 2H₂O → Ca(OH)₂ + C₂H₂↑ (ethyne/acetylene)

When heated with atmospheric nitrogen at 1100°C, CaC₂ forms calcium cyanamide (CaNCN) — an important slow-release nitrogenous fertiliser.

CaC₂ + N₂ →^1100°C CaNCN + C

Be₂C is unusual — it is a methanide carbide that reacts with water to give methane (not ethyne):

Be₂C + 4H₂O → 2Be(OH)₂ + CH₄↑

Hydrides

Ca, Sr, and Ba form ionic hydrides (MH₂) containing the hydride ion H⁻. BeH₂ and MgH₂ are covalent and polymeric. The polymeric (BeH₂)_n involves remarkable three-centre two-electron bonds (banana bonds) — a concept central to cluster chemistry.

CaH₂ + 2H₂O → Ca(OH)₂ + 2H₂↑

16. Grignard Reagents — Magnesium's Greatest Gift to Organic Chemistry

This section bridges inorganic and organic chemistry and is extremely important for JEE Advanced, IIT-JAM, and GATE. Victor Grignard, a Frenchman, won the Nobel Prize in Chemistry in 1912 for discovering organomagnesium compounds now bearing his name.

Preparation of Grignard Reagents

An alkyl or aryl halide (R–X, where X = Cl, Br, or I) is added slowly to magnesium turnings in absolutely dry diethyl ether. Water and air must be rigorously excluded:

Mg + RBr →^{dry ether} RMgBr (Grignard reagent)

Example: Mg + C₂H₅Br →^Et₂O C₂H₅MgBr (ethyl magnesium bromide)

Key points: Iodides are most reactive; chlorides are least reactive. Alkyl Grignard reagents are more reactive than aryl ones. Water rapidly destroys Grignard reagents by hydrolysis. The Grignard reagent is not isolated — it is made and used directly.

Key Reactions of Grignard Reagents

Figure 5: Reaction map of Grignard reagents — versatile in synthesising carboxylic acids, primary, secondary, and tertiary alcohols, and organosilicone compounds.

RMgBr + CO₂ →^{+ H⁺} R–COOH (carboxylic acid)

RMgBr + HCHO →^{+ H⁺} RCH₂OH (primary alcohol)

RMgBr + R'CHO →^{+ H⁺} R'(R)CHOH (secondary alcohol)

RMgBr + R'₂C=O →^{+ H⁺} R'₂(R)COH (tertiary alcohol)

RMgBr + I₂ → RI + MgBrI

RMgBr + SiCl₄ → RSiCl₃, R₂SiCl₂, R₃SiCl, R₄Si (silicones precursors)

2RMgBr + 2H₂O → 2RH + Mg(OH)₂ + MgBr₂ (hydrolysis)

17. Biological Roles of Mg²⁺ and Ca²⁺

Chemistry is not just about test tubes — it is the language of life. Mg²⁺ and Ca²⁺ are both essential elements in the human body, playing roles you simply cannot overlook:

Mg²⁺: Concentrated inside animal cells (like K⁺). Forms a complex with ATP — essential for all energy-releasing reactions. Required for impulse transmission along nerve fibres. At the centre of chlorophyll, the green pigment that drives photosynthesis and ultimately sustains all life on Earth.
Ca²⁺: Concentrated in body fluids outside cells (like Na⁺). Present in bones and teeth as apatite Ca₅(PO₄)₃(OH) and fluorapatite [3Ca₃(PO₄)₂·CaF₂]. Triggers muscle contraction (including heartbeat). Essential for blood clotting. Required for transmission of nerve impulses.

Figure 6: Simplified representation of the Mg²⁺ coordination environment in chlorophyll — the green pigment of life. Mg sits at the centre of a flat porphyrin ring, coordinated to four nitrogen atoms.

18. Extraction of Group 2 Metals

These metals cannot be extracted by simple chemical reduction because they are themselves strong reducing agents and react with carbon to form carbides. Aqueous electrolysis fails because they react violently with water. The main methods are:

Magnesium — Two Industrial Routes

Dow Seawater Process: Sea water contains ~0.13% Mg²⁺. Slaked lime is added to precipitate Mg(OH)₂, which is treated with HCl to give MgCl₂, then electrolysed.

Ca(OH)₂ + MgCl₂ → Mg(OH)₂↓ + CaCl₂

Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O

MgCl₂ →^electrolysis Mg + Cl₂

Pidgeon Process: Calcined dolomite [CaO·MgO] is reduced with ferrosilicon at 1150°C under reduced pressure:

[CaCO₃·MgCO₃] →^heat CaO·MgO + 2CO₂

2(CaO·MgO) + Si + Fe →^1150°C 2Mg↑ + Ca₂SiO₄ + Fe

Beryllium

BeO is extracted from beryl, converted to BeCl₂, then electrolysed. Alternatively, BeF₂ is reduced with magnesium metal.

BeF₂ + Mg →^heat Be + MgF₂

19. Exam Tips, Important Tricks and Highlights

🎯 Top 10 Most Asked Points — JEE/NEET/GATE

Thermal stability of carbonates and sulphates increases down the group (Be → Ba). BeCO₃ decomposes below 100°C; BaCO₃ at 1360°C.
Solubility of sulphates decreases down the group. BaSO₄ is virtually insoluble.
Solubility of hydroxides increases down the group. Ba(OH)₂ is strongly basic.
Be forms covalent compounds; all others form predominantly ionic compounds.
BeCl₂ in gas phase is linear (sp hybrid); in solid state it is a chain polymer with Cl bridges.
Only Ba forms peroxide (BaO₂) among Group 2 metals under ordinary conditions.
Be is amphoteric; Mg, Ca, Sr, Ba are purely basic.
Beryllium has a diagonal relationship with Aluminium.
Grignard reagent (RMgX) is prepared in dry diethyl ether; iodides react fastest.
Mg²⁺ is at the centre of chlorophyll; Ca²⁺ triggers muscle contraction and is in bones.

🧠 Solubility Trends Shortcut Think of two opposite rivers: River 1 (flows downhill = decreasing): Sulphates, Carbonates, Chromates, Oxalates. River 2 (flows uphill = increasing): Hydroxides, Fluorides. BaSO₄ → insoluble (used in X-ray). Be(OH)₂ → barely soluble; Ba(OH)₂ → quite soluble.

⚠️ Common Mistakes to Avoid

Do not confuse "temporary hardness" (due to bicarbonates, removed by boiling) with "permanent hardness" (due to sulphates/chlorides, removed by ion exchange or Na₂CO₃).
Be(OH)₂ is amphoteric — it is NOT a strong base like Ba(OH)₂. This trips up many students.
Mg does NOT react with cold water (oxide layer protects it), but reacts with hot water/steam.
In Grignard reagent preparation, even traces of moisture destroy the reagent. The solvent must be perfectly dry.
BaO₂ is a peroxide, not a superoxide. Only K, Rb, Cs form superoxides (MO₂) in Group 1.
Plaster of Paris has formula CaSO₄·½H₂O — the half-formula is exact and important.

20. Quick Revision Summary

Topic	Key Trend / Rule
Atomic size	Increases Be → Ra; smaller than Group 1
Ionisation energy	Decreases Be → Ra; IE₃ always very high
Electronegativity	Decreases Be(1.5) → Ba(0.9)
Hydration energy	Decreases Be → Ba (larger ion, weaker attraction)
Melting points	Higher than Group 1 (2 valence e⁻ in metallic bond)
Carbonate stability	Increases Be → Ba
Sulphate solubility	Decreases Be → Ba
Hydroxide solubility	Increases Be → Ba
Basic strength of hydroxides	Increases Be → Ba
Fluoride solubility	Increases Be → Ba
Reactivity with water	Increases Be → Ba (Be/Mg least reactive)
Covalent character	Most in Be compounds (highest IE, smallest size)
Peroxide formation	Only Ba forms BaO₂ under ordinary conditions
Diagonal relationship	Be ↔ Al (Group 13)

References: J.D. Lee, Concise Inorganic Chemistry, 5th Ed., Blackwell Science. IUPAC recommendations for nomenclature and standard electrode potentials. Crystal structure data from Inorganic Crystal Structure Database (ICSD). All reaction equations are verified against standard inorganic chemistry literature and IUPAC conventions.