Group 2: properties, reactions of elements and compounds, trends

Group 2 (Alkaline‑Earth Metals) – Cambridge AS & A‑Level Chemistry (9701) Syllabus

The Group 2 elements (Be, Mg, Ca, Sr, Ba, Ra) are metallic, form only a +2 oxidation state, and show clear, systematic trends in physical and chemical properties down the group. The notes below are organised to meet each specific syllabus requirement.

1. Physical‑property trends

*Values for radium are approximate (radioactive, limited experimental data).

• Atomic radius: Increases down the group as each successive element adds a new electron shell.
• First ionisation energy: Decreases down the group because the valence electrons are farther from the nucleus and are increasingly shielded.
• Melting point & density:
  • From Be to Ca the melting point falls – the larger ions have weaker metallic‑bond strength (lower lattice energy) and the crystal structures become less efficiently packed.
  • From Ca to Ba (and Ra) the melting point rises again as the heavier ions increase the mass of the lattice and the metallic bonding becomes more delocalised, giving higher density and a slight recovery in melting point.
• Electrical conductivity: Low compared with transition metals but rises down the group because the larger, more polarisable metal atoms allow the conduction electrons to move more freely.
• Lattice energy: Decreases down the group (larger ionic radii → larger inter‑ionic distances). This underpins the trends in melting point, density and the increasing basicity of oxides/hydroxides.

2. Chemical‑property trends

2.1 Reactivity

• With oxygen: All Group 2 metals react readily with O₂ to give the oxide M O. The rate of reaction increases down the group because the metals become more metallic (lower ionisation energy).
• With water:
  • Be – no reaction with cold water; reacts with steam (see 2.4).
  • Mg – reacts very slowly with cold water; reacts readily with hot water/steam.
  • Ca, Sr, Ba, Ra – react vigorously with cold water, producing a strong alkaline solution and hydrogen gas.
• With acids (e.g., HCl): All dissolve, liberating H₂ gas. Reaction rate follows the same trend as water reactivity.
• With halogens: Direct combination at room temperature (or with gentle heating) to give the halide M X₂ (X = F, Cl, Br, I).

2.2 Basicity of oxides and hydroxides

Corresponding hydroxides give very similar pH values; the trend is the same – basicity increases down the group because the oxide/hydroxide ion becomes larger and more polarisable, stabilising the negative charge.

2.3 Solubility trends (25 °C, water)

• Hydroxides: Solubility increases down the group because the lattice energy of the solid falls faster than the hydration energy of the ion.
• Carbonates: All are sparingly soluble; solubility decreases down the group as lattice energy increases.
• Sulfates: BaSO₄ and RaSO₄ are the notable exceptions – the large cation produces a very high lattice energy that outweighs hydration.
• Halides: BeCl₂ is covalent and polymeric, giving only limited solubility; the others are ionic and fully soluble.
• Nitrates: All Group 2 nitrates are highly soluble – a useful point for precipitation‑type questions.

2.4 Thermal decomposition of carbonates (syllabus requirement)

When heated, each Group 2 carbonate decomposes to the corresponding oxide and carbon dioxide:

 MCO₃ (s) →[Δ] MO (s) + CO₂ (g) 

The temperature required decreases down the group (e.g., MgCO₃ ≈ 350 °C, CaCO₃ ≈ 825 °C, SrCO₃ ≈ 1000 °C, BaCO₃ ≈ 1100 °C). This trend follows the decreasing lattice energy of the carbonate.

3. General reactions of Group 2 metals (M)

1. Combustion (oxygen): M (s) + ½ O₂ (g) → MO (s)
2. Reaction with water (cold): M (s) + 2 H₂O (l) → M(OH)₂ (aq) + H₂ (g)
   • Applicable to Ca, Sr, Ba, Ra.
3. Reaction with hot water/steam: M (s) + H₂O (g) → M(OH)₂ (aq) + H₂ (g)
   • Mg – hot water; Be – steam only (produces BeO + H₂).
4. Reaction with acids (e.g., HCl): M (s) + 2 HCl (aq) → MCl₂ (aq) + H₂ (g)
5. Reaction with halogens: M (s) + X₂ (g) → MX₂ (s) (X = F, Cl, Br, I)
6. Thermal decomposition of carbonates: MCO₃ (s) →[Δ] MO (s) + CO₂ (g)
7. Formation of hydroxides from oxides: MO (s) + H₂O (l) → M(OH)₂ (aq)

4. Representative examples (for exam answers)

• Calcium + cold water: Ca (s) + 2 H₂O (l) → Ca(OH)₂ (aq) + H₂ (g)
• Magnesium + oxygen (flame test): 2 Mg (s) + O₂ (g) → 2 MgO (s)
• Strontium oxide + water: SrO (s) + H₂O (l) → Sr(OH)₂ (aq)
• Barium carbonate decomposition: BaCO₃ (s) →[Δ] BaO (s) + CO₂ (g)
• Magnesium + hydrochloric acid: Mg (s) + 2 HCl (aq) → MgCl₂ (aq) + H₂ (g)
• Calcium + chlorine gas: Ca (s) + Cl₂ (g) → CaCl₂ (s)
• Beryllium + steam (exception): Be (s) + H₂O (g) → BeO (s) + H₂ (g)

5. Exam‑style checklist

• Start answers with the generic equation for the whole group, then give a specific metal if the question asks.
• When discussing trends, always link to ionisation energy, atomic radius and lattice energy.
• Remember the **Be exception** – no reaction with cold water; reacts with steam to give BeO.
• For solubility questions, use the following quick‑reference:
  • Hydroxides: solubility ↑ down the group.
  • Carbonates: sparingly soluble, solubility ↓ down the group.
  • Sulfates: all soluble except BaSO₄ (and RaSO₄).
  • Halides: all soluble except polymeric BeCl₂.
  • Nitrates: all soluble.
• State pH values (or “strong base”) for 0.1 M oxide/hydroxide solutions to earn marks for the basicity outcome.
• Explain the melting‑point dip (Be → Ca) and subsequent rise (Ca → Ba) in terms of metallic‑bond strength, packing efficiency and lattice‑energy changes.
• Balance all equations; remember the metal always forms a +2 ion, giving formulas MX₂, MCO₃, MSO₄, etc.