Group 2 Compounds

Group 2 compounds — particularly the oxides, hydroxides, carbonates, and nitrates — display systematic trends that are directly linked to the size and charge density of the metal cation. Understanding these trends is essential for Edexcel A-Level questions on thermal stability, acid-base behaviour, and real-world applications.

Thermal Decomposition of Carbonates

Group 2 carbonates decompose on heating to form the metal oxide and carbon dioxide:

MCO₃(s) → MO(s) + CO₂(g)

The key trend is that thermal stability increases down the group. MgCO₃ decomposes at roughly 350 °C, while BaCO₃ requires temperatures above 1300 °C.

Why Does Stability Increase?

The explanation lies in the polarising power of the metal cation. Smaller cations (like Mg²⁺) have a high charge density, meaning their positive charge is concentrated in a small volume. This high charge density distorts (polarises) the electron cloud of the large carbonate ion (CO₃²⁻), weakening the C–O bonds within it and making decomposition easier.

As you descend the group, the cations become larger, their charge density decreases, and they polarise the carbonate ion less effectively. The C–O bonds remain stronger, and a higher temperature is needed to decompose the carbonate.

Carbonate	Decomposition Temp / °C	Cation Radius / pm	Charge Density / C mm⁻³
MgCO₃	~350	72	High
CaCO₃	~840	100	Moderate
SrCO₃	~1100	118	Lower
BaCO₃	~1360	135	Lowest

flowchart LR
    subgraph "Thermal Stability of Carbonates"
        A["MgCO₃<br>~350 °C<br>Least stable"] -->|"Increasing stability"| B["CaCO₃<br>~840 °C"] -->|"Increasing stability"| C["SrCO₃<br>~1100 °C"] -->|"Increasing stability"| D["BaCO₃<br>~1360 °C<br>Most stable"]
    end
    E["Explanation: Cation size increases → charge density decreases → less polarisation of CO₃²⁻"]

Common exam mistake: Students often say "the bond between the metal and the carbonate breaks more easily." This is wrong. It is the bonds within the carbonate ion (C–O bonds) that are weakened by polarisation. The metal cation distorts the anion's electron cloud, not its own bond to the anion. Examiners specifically penalise this error.

Thermal Decomposition of Nitrates

Group 2 nitrates decompose on heating, but the products differ from the carbonates:

2M(NO₃)₂(s) → 2MO(s) + 4NO₂(g) + O₂(g)

Brown fumes of nitrogen dioxide (NO₂) are observed, along with oxygen gas. The same thermal stability trend applies: Mg(NO₃)₂ decomposes at a lower temperature than Ba(NO₃)₂, again because the smaller Mg²⁺ ion polarises the nitrate ion more effectively.

The nitrate ion (NO₃⁻) is also a large anion with a delocalised electron cloud that can be distorted by the cation's electric field.

Contrast with Group 1 nitrates: Group 1 nitrates (except lithium) decompose to the nitrite and oxygen only: 2NaNO₃ → 2NaNO₂ + O₂. This is because Group 1 cations are M⁺ (lower charge) and therefore polarise the nitrate ion less, so decomposition is less complete. Lithium nitrate decomposes like a Group 2 nitrate (to Li₂O, NO₂, and O₂) because Li⁺ has an exceptionally high charge density due to its tiny ionic radius.

Oxides and Their Reactions with Water

Group 2 oxides react with water to form hydroxides:

MO(s) + H₂O(l) → M(OH)₂(aq or s)

Reaction	Observation	pH of resulting solution
MgO + H₂O → Mg(OH)₂	Slow; weakly alkaline suspension	~10
CaO + H₂O → Ca(OH)₂	Exothermic; slaked lime forms	~12.4
SrO + H₂O → Sr(OH)₂	Exothermic; more alkaline	~13
BaO + H₂O → Ba(OH)₂	Exothermic; strongly alkaline	~13.5

Group 2 oxides are basic oxides. They react with acids to form a salt and water:

MgO(s) + 2HCl(aq) → MgCl₂(aq) + H₂O(l)

Hydroxides as Bases

Group 2 hydroxides are bases. They neutralise acids:

M(OH)₂ + 2HCl → MCl₂ + 2H₂O

Because solubility of hydroxides increases down the group, Ba(OH)₂ is a stronger base in solution than Mg(OH)₂. Solutions of Ba(OH)₂ have a higher pH than saturated Mg(OH)₂ solutions.

Uses of Group 2 Compounds

Several Group 2 compounds have important practical applications that appear regularly in exam questions:

Mg(OH)₂ — Antacids

Magnesium hydroxide is used in antacid tablets and suspensions (such as milk of magnesia) to neutralise excess stomach acid (HCl):

Mg(OH)₂(s) + 2HCl(aq) → MgCl₂(aq) + 2H₂O(l)

It is chosen because it is insoluble enough to be safe to ingest but reacts with acid effectively. It is a weak base that neutralises acid without making the stomach dangerously alkaline.

Ca(OH)₂ — Agriculture and Water Treatment

Calcium hydroxide (slaked lime) is used by farmers to neutralise acidic soils:

Ca(OH)₂(s) + H₂SO₄(aq) → CaSO₄(aq) + 2H₂O(l)

Acidic soils (low pH) inhibit plant growth. Adding Ca(OH)₂ raises the pH to a more suitable range (ideally pH 6–7 for most crops).

Ca(OH)₂ is also used in water treatment to raise the pH of acidic water supplies and to remove temporary hardness:

Ca(OH)₂(aq) + Ca(HCO₃)₂(aq) → 2CaCO₃(s) + 2H₂O(l)

CaO — Water Treatment

Calcium oxide (quicklime) reacts exothermically with water:

CaO(s) + H₂O(l) → Ca(OH)₂(aq)

CaO is sometimes preferred over Ca(OH)₂ because it is lighter and cheaper to transport — the water at the treatment site provides the H₂O needed to form the hydroxide in situ.

BaSO₄ — Barium Meal

Barium sulfate is used in medical imaging as a "barium meal." Patients swallow a suspension of BaSO₄ before an X-ray of the digestive system. Barium is dense and absorbs X-rays well, providing contrast in the image. Despite barium ions being toxic, BaSO₄ is safe because it is completely insoluble — no Ba²⁺ ions enter the bloodstream.

Linking Thermal Stability to Polarisation

The concept of polarisation (charge density of the cation distorting the anion) is a recurring theme. To score full marks on exam questions, you should be able to:

State the trend (stability increases down the group)
Explain it in terms of cation size and charge density
Link charge density to the polarisation of the anion
Explain that greater polarisation weakens bonds within the anion, making decomposition easier

Worked Example: Exam-Style 6-Mark Answer

Question: Explain why MgCO₃ has a lower thermal decomposition temperature than BaCO₃. [6 marks]

Model answer:

Mg²⁺ has a smaller ionic radius than Ba²⁺ (72 pm vs 135 pm) [1 mark]
Both ions have a +2 charge, so Mg²⁺ has a higher charge density [1 mark]
Mg²⁺ polarises the carbonate ion (CO₃²⁻) more strongly [1 mark]
This distorts the electron cloud of the CO₃²⁻ ion [1 mark]
Weakening the C–O bonds within the carbonate ion [1 mark]
Less energy is required to break these weakened bonds, so decomposition occurs at a lower temperature [1 mark]

Examiner tip: Never write "weaker bond between metal and carbonate" — this is a common error that will lose marks.

Summary

Group 2 carbonates and nitrates become more thermally stable down the group because the increasing cation size reduces polarisation of the anion. Oxides react with water to form hydroxides, which act as bases. Practical applications include Mg(OH)₂ in antacids, Ca(OH)₂ in agriculture, CaO in water treatment, and BaSO₄ in barium meals. Always explain thermal stability trends using charge density and polarisation of the anion's internal bonds.

A-Level Deep Dive: Group 2 Compounds

Spec mapping

Edexcel 9CH0 specification Topic 4 — Inorganic Chemistry and the Periodic Table, sub-topic 4.1 covers trends in the solubility of Group 2 hydroxides and sulfates, and trends in the thermal stability of Group 2 carbonates and nitrates (refer to the official specification document for exact wording). The material is examined directly in Paper 1 (9CH0/01) and synoptically in Paper 3 (9CH0/03) through CP3 observations and qualitative-analysis sequences. Underpinning concepts come from Topic 13 (Energetics II: Born-Haber and lattice energy), Topic 11 (Equilibrium I: solubility and Δ_solH) and Topic 1 (Atomic Structure) for the cation-size argument that drives every trend in this sub-topic.

Worked example with full mark scheme

Question (8 marks):

(a) Predict and explain the trend in solubility of Group 2 sulfates (MgSO₄ → BaSO₄), referring to lattice energy and hydration enthalpy. (5)

(b) Explain why MgCO₃ decomposes at a much lower temperature than BaCO₃. (3)

Solution with mark scheme:

(a) Step 1 — state the trend.

Solubility of MSO₄ decreases down the group: MgSO₄ ≈ 35 g/100 g H₂O at 25 °C is highly soluble; CaSO₄ ≈ 0.21 g; SrSO₄ ≈ 0.013 g; BaSO₄ ≈ 0.00024 g (a six-order-of-magnitude fall).

B1 — correct trend stated quantitatively.

Step 2 — set up the energetic competition.

The enthalpy of solution Δ_solH is approximately:

$\Delta_{sol}H \approx -\Delta_{lat}H + (\Delta_{hyd}H_{cation} + \Delta_{hyd}H_{anion})$

For dissolution to be favourable, hydration must compensate for breaking the lattice.

M1 — Hess-cycle setup recognising both lattice and hydration contributions.

Step 3 — analyse the changes down the group.

Both lattice energy and hydration enthalpy of M²⁺ become less negative as cation radius increases (1/r dependence in both Born-Landé and ion-dipole models). The crucial point: the SO₄²⁻ anion is large, so the lattice energy is dominated by anion size and changes only slowly down the group; in contrast, ΔhydH(M²⁺) depends on the cation alone and falls rapidly.

M1 — recognition that lattice energy changes less than hydration enthalpy because the large anion dominates ΔlatH.

A1 — therefore the magnitude by which hydration compensates for the lattice falls faster than the lattice itself shrinks; net Δ_solH becomes less exothermic (or more endothermic), so solubility decreases.

A1 — final synthesis: BaSO₄ is essentially insoluble because hydration of Ba²⁺ cannot pay back the BaSO₄ lattice.

(b) Step 1 — polarising power of M²⁺.

Charge density of Mg²⁺ (small cation) is far higher than Ba²⁺. Mg²⁺ polarises the electron cloud of CO₃²⁻, distorting the C–O bonds and weakening one of them.

M1 — polarising-power argument.

Step 2 — energetic restatement.

The decomposition MCO₃(s) → MO(s) + CO₂(g) is favoured when ΔlatH(MO) − ΔlatH(MCO₃) is large. Because O²⁻ is a smaller anion than CO₃²⁻, the lattice energy gain on forming MO is much larger when M²⁺ is small (Mg²⁺): MgO is exceptionally stable.

M1 — lattice-energy rationale linking small M²⁺ + small O²⁻ to high MO lattice stability.

A1 — therefore MgCO₃ → MgO + CO₂ has a more negative ΔH at lower T; decomposition temperature is ~350 °C for MgCO₃ vs ~1300 °C for BaCO₃.

Total: 8 marks (M3 A3 B1 + 1 calc).

Specimen question modelled on the Edexcel 9CH0 Paper 1 format

Question (6 marks): Indigestion remedies often contain magnesium hydroxide.

(a) Write the equation for the reaction of Mg(OH)₂ with HCl. (1)

(b) State and explain whether you would expect Sr(OH)₂ or Mg(OH)₂ to be more soluble in water. (3)

(c) Suggest, with reasoning, why Ba(OH)₂ is not used in indigestion remedies despite being even more soluble. (2)

Mark scheme decomposition by AO:

Mark	AO	Awarded for
(a) B1	AO1	Mg(OH)₂(s) + 2HCl(aq) → MgCl₂(aq) + 2H₂O(l) — correctly balanced + state symbols
(b) B1	AO2	Sr(OH)₂ is more soluble
(b) M1	AO2	Hydroxide solubility increases down Group 2
(b) A1	AO3	Lattice energy and hydration enthalpy of M²⁺ both fall, but for the small OH⁻ anion lattice energy falls faster than hydration → ΔsolH more exothermic for larger M²⁺
(c) M1	AO1	Ba²⁺ is toxic / poisonous
(c) A1	AO3	Ba(OH)₂ is also strongly alkaline (corrosive at the higher concentrations its solubility allows) — both safety and toxicity preclude pharmaceutical use

Total: 6 marks split AO1 = 2, AO2 = 2, AO3 = 2. Group 2 compound questions blend factual recall with a thermodynamic justification, so AO3 reasoning marks are commonly available even on apparently short-answer questions.