Group 2: The Alkaline Earth Metals

Group 2 of the periodic table contains the alkaline earth metals: beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), and barium (Ba). At A-Level, the focus is on magnesium to barium, since beryllium behaves anomalously due to its very small atomic radius and high charge density. These elements share a common electron configuration ending in s², meaning each atom has two electrons in its outermost shell. This accounts for the +2 oxidation state that dominates their chemistry.

Trends in Atomic Radius

As you descend Group 2, atomic radius increases. Each successive element has an additional electron shell, so the outermost electrons are further from the nucleus.

Element	Atomic Radius / pm	Electron Configuration	1st IE / kJ mol⁻¹	2nd IE / kJ mol⁻¹
Mg	160	1s² 2s² 2p⁶ 3s²	738	1451
Ca	197	[Ar] 4s²	590	1145
Sr	215	[Kr] 5s²	550	1064
Ba	222	[Xe] 6s²	503	965

The increase in radius has important consequences for reactivity, ionisation energy, and the properties of compounds.

Trends in Ionisation Energy

First ionisation energy decreases down Group 2. Although nuclear charge increases (more protons), the effect is outweighed by increased shielding from inner electron shells and greater distance of the outer electrons from the nucleus. The outermost electrons are easier to remove in barium than in magnesium.

The second ionisation energy also decreases down the group for the same reasons. Since Group 2 elements form M²⁺ ions by losing both s² electrons, the decreasing ionisation energies explain the increasing reactivity down the group.

The jump from the 2nd to the 3rd ionisation energy is enormous for every Group 2 element (for Mg, the 3rd IE is 7733 kJ mol⁻¹ compared to 1451 kJ mol⁻¹ for the 2nd). This is because the third electron would be removed from an inner shell, which is much closer to the nucleus and far more strongly held. This is why Group 2 elements never form M³⁺ ions.

Melting Points

Melting points do not follow a simple trend in Group 2. The overall tendency is for melting point to decrease as the metallic radius increases and the bond strength per unit volume decreases, but changes in crystal packing cause irregularities.

Element	Melting Point / °C	Crystal Structure
Be	1287	HCP
Mg	650	HCP
Ca	842	FCC
Sr	777	FCC
Ba	727	BCC

The irregularity at calcium (higher than magnesium) arises because calcium adopts a face-centred cubic structure rather than hexagonal close packed, which affects the efficiency of metallic bonding.

Reactions with Water

The Group 2 metals react with water to form a metal hydroxide and hydrogen gas:

M(s) + 2H₂O(l) → M(OH)₂(aq) + H₂(g)

Reactivity increases down the group because ionisation energies decrease, making it easier for atoms to lose their two outer electrons.

Magnesium reacts very slowly with cold water, forming a thin layer of Mg(OH)₂ on the surface. It reacts vigorously with steam: Mg(s) + H₂O(g) → MgO(s) + H₂(g). Note the different product with steam (oxide, not hydroxide).
Calcium reacts steadily with cold water, producing bubbles of hydrogen and a white cloudy suspension of Ca(OH)₂ (limewater). The solution is mildly alkaline (pH ≈ 12.4 for saturated limewater).
Strontium reacts more vigorously than calcium with cold water, producing faster bubbling and a more alkaline solution.
Barium reacts vigorously with cold water, quickly producing hydrogen gas and dissolving to form Ba(OH)₂. The solution is strongly alkaline.

Common exam mistake: Students sometimes write Mg(OH)₂ as the product when magnesium reacts with steam. With steam, the product is MgO (not the hydroxide), because the high temperature decomposes any hydroxide formed.

Flame Colours

When Group 2 metal compounds are heated in a flame, electrons are promoted to higher energy levels and then fall back, emitting light of characteristic wavelengths.

Element	Flame Colour	Approximate Wavelength / nm
Mg	No distinctive colour (white/bright)	—
Ca	Brick red / orange-red	622 (orange-red)
Sr	Crimson / deep red	675 (deep red)
Ba	Pale green / apple green	524 (green)

These flame colours are used as a diagnostic test to identify Group 2 metal ions. It is important to remember that the colours arise from electron transitions, not from the ion itself being coloured. The energy gap between the excited state and the ground state determines the wavelength of light emitted.

flowchart TD
    A["Unknown Group 2 compound"] --> B{"Flame test colour?"}
    B -->|"Brick red / orange-red"| C["Ca²⁺"]
    B -->|"Crimson / deep red"| D["Sr²⁺"]
    B -->|"Pale green / apple green"| E["Ba²⁺"]
    B -->|"No distinctive colour"| F["Mg²⁺ (use other tests)"]

Solubility Trends

Two important solubility trends must be learned for Group 2 compounds. These trends are based on the balance between lattice enthalpy and hydration enthalpy of the ions, and both decrease down the group, but at different rates.

Hydroxides: Solubility Increases Down the Group

Compound	Solubility at 25 °C / mol dm⁻³	Description
Mg(OH)₂	2.0 × 10⁻⁴	Almost insoluble; forms a suspension
Ca(OH)₂	1.5 × 10⁻²	Slightly soluble; limewater is a saturated solution
Sr(OH)₂	3.4 × 10⁻²	Moderately soluble
Ba(OH)₂	1.5 × 10⁻¹	Most soluble; strongly alkaline

This trend means that as you go down Group 2, the hydroxides dissolve more readily and produce more alkaline solutions. Ba(OH)₂ solution is strongly alkaline (pH ≈ 13–14 depending on concentration).

Sulfates: Solubility Decreases Down the Group

Compound	Solubility at 25 °C / mol dm⁻³	Description
MgSO₄	2.1	Very soluble (Epsom salts)
CaSO₄	4.7 × 10⁻²	Slightly soluble
SrSO₄	5.3 × 10⁻⁴	Sparingly soluble
BaSO₄	9.4 × 10⁻⁶	Insoluble

The insolubility of BaSO₄ is exploited in the test for sulfate ions: adding BaCl₂ solution to a sample acidified with HCl produces a white precipitate of BaSO₄ if sulfate ions are present.

Why Do These Trends Occur?

For both hydroxides and sulfates, lattice enthalpy decreases down the group (as the cation gets bigger, the ions are further apart and the lattice is less strongly held). Hydration enthalpy also decreases (less energy is released when larger ions are hydrated). The key is that for hydroxides, the lattice enthalpy decreases faster than hydration enthalpy, so dissolving becomes more favourable. For sulfates, the hydration enthalpy decreases faster than lattice enthalpy, so dissolving becomes less favourable.

flowchart LR
    subgraph "Hydroxide Solubility"
        A1["Mg(OH)₂ – Insoluble"] --> B1["Ca(OH)₂ – Slight"] --> C1["Sr(OH)₂ – Moderate"] --> D1["Ba(OH)₂ – Soluble"]
    end
    subgraph "Sulfate Solubility"
        A2["MgSO₄ – Soluble"] --> B2["CaSO₄ – Slight"] --> C2["SrSO₄ – Sparingly"] --> D2["BaSO₄ – Insoluble"]
    end

Worked Example: Identifying an Unknown Group 2 Compound

A student dissolves a white solid in water. The resulting solution is strongly alkaline (pH 13). They add dilute H₂SO₄ to a portion and observe a white precipitate.

Analysis:

Strongly alkaline → the hydroxide must be very soluble → likely Ba(OH)₂
White precipitate with H₂SO₄ → BaSO₄ (insoluble sulfate confirms Ba²⁺)
Confirmation: flame test would show pale green

Answer: The compound is barium hydroxide, Ba(OH)₂.

If the solution had been only weakly alkaline (pH ≈ 10–11) and no precipitate formed with dilute H₂SO₄, the compound could be Ca(OH)₂, since CaSO₄ is slightly soluble and may not produce a clear precipitate.

Summary

Group 2 shows clear periodic trends driven by increasing atomic radius and decreasing ionisation energy. Reactivity with water increases down the group, hydroxide solubility increases, and sulfate solubility decreases. Flame tests provide a quick way to identify individual Group 2 elements. Always include actual data values (ionisation energies, solubilities, decomposition temperatures) in your exam answers to demonstrate thorough knowledge.