Group 2 Trends and Reactions
Group 2 Trends and Reactions
Group 2 elements (Be, Mg, Ca, Sr, Ba) are the alkaline earth metals. They show clear trends in physical and chemical properties descending the group, driven by increasing atomic radius and decreasing ionisation energy.
Trends Down Group 2
Atomic radius increases — additional electron shells are added.
First ionisation energy decreases — the outer 2 electrons are further from the nucleus and more shielded, so less energy is needed to remove one.
Reactivity increases — it becomes easier to lose the two outer electrons, so reactions occur more readily.
Melting point — the general trend is a decrease down the group, though it is not perfectly regular. Metallic bonding weakens as the atoms get larger (delocalised electrons are further from the positive nuclei).
Electronegativity decreases — larger atoms attract bonding electrons less strongly.
Electron Configuration
All Group 2 elements have the configuration [noble gas] ns², where n is the period number. They form M²⁺ ions by losing both outer s electrons.
Reactions with Water
Group 2 metals react with water to form a hydroxide and hydrogen gas:
M(s) + 2H₂O(l) → M(OH)₂(aq) + H₂(g)
The reactivity increases down the group:
- Magnesium reacts very slowly with cold water (forms a thin oxide layer) but burns vigorously in steam:
Mg(s) + H₂O(g) → MgO(s) + H₂(g)
- Calcium reacts steadily with cold water, producing bubbles of hydrogen and a milky suspension of Ca(OH)₂
- Strontium reacts more vigorously than calcium
- Barium reacts vigorously with cold water
Beryllium does not react with water (the oxide layer is too stable and Be has very high ionisation energies).
Reactions with Oxygen
All Group 2 metals burn in oxygen to form white solid oxides:
2M(s) + O₂(g) → 2MO(s)
- Magnesium burns with an intense white flame (used in flares and fireworks)
- The oxides are ionic with the general formula MO, containing M²⁺ and O²⁻ ions
- Reactivity increases down the group
Reactions with Dilute Acids
Group 2 metals react with dilute hydrochloric acid:
M(s) + 2HCl(aq) → MCl₂(aq) + H₂(g)
And with dilute sulfuric acid:
M(s) + H₂SO₄(aq) → MSO₄(aq) + H₂(g)
Reactivity increases down the group. However, the reaction of barium with sulfuric acid appears to stop quickly because BaSO₄ is insoluble — it coats the surface of the metal, preventing further reaction.
Solubility Trends of Group 2 Compounds
Two key trends to learn:
Hydroxides — solubility INCREASES down the group:
- Mg(OH)₂ is sparingly soluble (milk of magnesia, used as an antacid)
- Ca(OH)₂ is slightly soluble (limewater — used to test for CO₂)
- Sr(OH)₂ is moderately soluble
- Ba(OH)₂ is soluble
This means the pH of the saturated hydroxide solutions increases down the group (more OH⁻ ions dissolve).
Sulfates — solubility DECREASES down the group:
- MgSO₄ is soluble (Epsom salts)
- CaSO₄ is slightly soluble
- SrSO₄ is sparingly soluble
- BaSO₄ is insoluble (used in barium meals for X-ray imaging of the gut — it is opaque to X-rays and its insolubility means it is not absorbed, so it is non-toxic despite barium ions being poisonous)
Thermal Decomposition of Group 2 Carbonates and Nitrates
Carbonates decompose on heating:
MCO₃(s) → MO(s) + CO₂(g)
The temperature required increases down the group:
- MgCO₃ decomposes easily (around 540 °C)
- CaCO₃ at ~840 °C
- BaCO₃ at ~1360 °C
Explanation: Larger M²⁺ ions (down the group) are less polarising — they distort the electron cloud of the CO₃²⁻ ion less, making it harder to break down the carbonate. More thermal energy is needed.
Nitrates decompose similarly:
2M(NO₃)₂(s) → 2MO(s) + 4NO₂(g) + O₂(g)
Again, thermal stability increases down the group for the same polarisation reason.
Uses of Group 2 Compounds
| Compound | Use | Reason |
|---|---|---|
| Mg(OH)₂ | Antacid (milk of magnesia) | Neutralises excess stomach acid |
| Ca(OH)₂ | Neutralising acidic soil | Cheap, slightly soluble base |
| Ca(OH)₂ | Limewater (CO₂ test) | Ca(OH)₂ + CO₂ → CaCO₃ (milky) |
| CaCO₃ | Cement and building | Thermal decomposition gives CaO |
| BaSO₄ | Barium meal (X-ray contrast) | Insoluble, opaque to X-rays |
Flame Colours
Group 2 metals produce characteristic flame colours when their compounds are heated:
- Ca: brick red / orange-red
- Sr: crimson red
- Ba: pale green
(Mg and Be do not produce visible flame colours in the same way.)
These colours arise because heat promotes electrons to higher energy levels; as they fall back, they emit photons of specific wavelengths.
Exam Tips
- Learn the two solubility trends (hydroxides UP, sulfates DOWN) — these are examined very frequently
- When explaining thermal stability of carbonates, always link to the polarising power of the cation (charge density = charge / radius)
- Write balanced equations including state symbols for all reactions