Group 2 Compounds and Uses

Spec Mapping — OCR H432 Module 3.1.2 — Group 2, covering the reactions of Group 2 oxides with water, the trends in solubility of Group 2 hydroxides (increasing down the group) and Group 2 sulfates (decreasing down the group), and the industrial / medicinal uses of selected Group 2 compounds: Mg(OH)₂ as a mild antacid, Ca(OH)₂ for liming acidic soils, CaCO₃ for flue-gas desulfurisation, and BaSO₄ as a contrast medium in barium meals (refer to the official OCR H432 specification document for exact wording).

The compounds of Group 2 — oxides, hydroxides, carbonates and sulfates — are produced industrially on a vast scale and find applications ranging from agriculture (soil liming with Ca(OH)₂) to power-station emissions control (CaCO₃ in flue-gas desulfurisation) to medical imaging (BaSO₄ for gastrointestinal X-rays). At the heart of OCR Module 3.1.2 are two solubility trends that move in opposite directions down the group: Group 2 hydroxides become more soluble down the group, while Group 2 sulfates become less soluble. The directions look counter-intuitive at first sight but follow from competing trends in lattice enthalpy and hydration enthalpy (Year 13 energetics fills in the thermodynamic detail). For OCR AS, you need to recall the direction of each trend, write the balanced equations for the relevant reactions, and link the solubility behaviour to the specific industrial or medicinal application — particularly the qualitative test for sulfate ions (Ba²⁺ + SO₄²⁻ → BaSO₄ white precipitate), which feeds into PAG 4 (qualitative analysis) directly. This lesson develops that fluency.

Key Definitions:

• Solubility — the maximum mass (or moles) of solute that dissolves in 100 g of water at a stated temperature; reported here in mol per 100 g H₂O at 25 °C.
• Slaking of lime — the exothermic reaction of CaO(s) with water to form Ca(OH)₂.
• Flue-gas desulfurisation (FGD) — the industrial process of removing SO₂ from power-station emissions using a calcium-containing alkaline reagent.
• Barium meal — a suspension of insoluble BaSO₄ swallowed by a patient to provide X-ray contrast in the gastrointestinal tract.

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Group 2 oxides with water

Group 2 oxides (MO) are ionic solids containing M²⁺ and O²⁻ ions. They react with water to form the corresponding metal hydroxides M(OH)₂:

$\text{MO(s)} + \text{H}_2\text{O(l)} \rightarrow \text{M(OH)}_2 \text{(aq or s)}$MO(s)+H2​O(l)→M(OH)2​(aq or s)

Oxide	Reaction with water	Solubility of product	Notes
MgO	Slow, partial	Mg(OH)₂ sparingly soluble (s)	Reaction sluggish at room T
CaO	Exothermic — "slaking of lime"	Ca(OH)₂ slightly soluble ("limewater")	Generates steam if mass of water is small
SrO	Vigorous	Sr(OH)₂ more soluble (aq)
BaO	Very vigorous	Ba(OH)₂ readily soluble (aq)

Worked equations:

$\text{CaO(s)} + \text{H}_2\text{O(l)} \rightarrow \text{Ca(OH)}_2\text{(aq)} \quad \Delta H \text{ exothermic}$CaO(s)+H2​O(l)→Ca(OH)2​(aq)ΔH exothermic

$\text{BaO(s)} + \text{H}_2\text{O(l)} \rightarrow \text{Ba(OH)}_2\text{(aq)}$BaO(s)+H2​O(l)→Ba(OH)2​(aq)

The resulting solutions are alkaline because OH⁻(aq) ions are released. Going down the group, the oxides react more vigorously and the resulting hydroxides are more soluble, so the OH⁻ concentration in saturated solution rises and the pH rises:

Hydroxide	Approx. pH of saturated solution
Mg(OH)₂	~10
Ca(OH)₂	~12
Sr(OH)₂	~13
Ba(OH)₂	~13–14

A saturated solution of Ca(OH)₂ is the classic limewater used in school labs to test for CO₂ gas: CO₂ bubbled through limewater forms a white CaCO₃ precipitate (limewater turns cloudy).

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Solubility trend 1: Group 2 hydroxides — increases down the group

Hydroxide	Solubility (mol / 100 g H₂O at 25 °C)	Common name / use
Mg(OH)₂	~0.00002	"Milk of magnesia" — antacid
Ca(OH)₂	~0.002	"Limewater" — slaked lime
Sr(OH)₂	~0.008	Less common laboratory base
Ba(OH)₂	~0.05	"Baryta water" — strong base

The numerical solubility rises by roughly three orders of magnitude from Mg(OH)₂ (essentially insoluble) to Ba(OH)₂ (fully soluble in modest amounts of water).

Group 2 sulfates MgSO₄ very soluble CaSO₄ slightly soluble SrSO₄ almost insoluble BaSO₄ insoluble Solubility ↓ down group

OCR does not require you to explain the thermodynamic reason for this trend in detail (that belongs to Year 13 energetics, where lattice enthalpy and hydration enthalpy compete). You need to:

• Recall the direction (hydroxide solubility ↑ down the group; sulfate solubility ↓ down the group).
• Apply it to the qualitative-test reasoning later in the course.
• Link to the industrial / medicinal applications (below).

A brief thermodynamic gloss for context: solubility is governed by the balance of lattice enthalpy (energy to break the ionic lattice; depends on inverse of (r⁺ + r⁻)) and hydration enthalpy (energy released when M²⁺ and X⁻ are hydrated; depends on inverse of cation/anion radius). For Group 2 hydroxides, the OH⁻ is small enough that the lattice enthalpy falls faster than the hydration enthalpy down the group → solubility increases. For Group 2 sulfates, SO₄²⁻ is much larger, so the hydration-enthalpy fall dominates the lattice-enthalpy fall → solubility decreases. (Year 13 will revisit this with full Born–Haber data.)

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Solubility trend 2: Group 2 sulfates — decreases down the group

Sulfate	Solubility (mol / 100 g H₂O)	Common name
MgSO₄	~0.3 (very soluble)	"Epsom salts"
CaSO₄	~0.0002 (sparingly soluble)	"Plaster of Paris" (gypsum dihydrate)
SrSO₄	~6 × 10⁻⁶	"Celestine"
BaSO₄	~1 × 10⁻⁸ (essentially insoluble)	"Barite" / "Barytes"

The solubility falls by roughly seven orders of magnitude from MgSO₄ to BaSO₄ — even larger than the hydroxide trend.

BaSO₄'s extreme insolubility has two famous applications:

1. Qualitative test for sulfate ions (PAG 4): add a few drops of acidified BaCl₂ solution to the unknown solution. A white precipitate of BaSO₄ confirms SO₄²⁻ is present:

   $\text{Ba}^{2+}\text{(aq)} + \text{SO}_4^{2-}\text{(aq)} \rightarrow \text{BaSO}_4\text{(s)}$Ba2+(aq)+SO42−​(aq)→BaSO4​(s)

   Acidification with dilute HCl removes any carbonate (which would also give a white precipitate of BaCO₃) — a fundamental rule revisited in the qualitative-analysis lesson.

2. Barium meals in medical imaging (see Uses below) — exploits BaSO₄'s insolubility in water and dilute stomach acid to immobilise toxic Ba²⁺.

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Key uses of Group 2 compounds

1. Mg(OH)₂ — "Milk of Magnesia" antacid

Mg(OH)₂ is a mild, sparingly soluble base. Because it is insoluble, the pH of the gastric environment is never raised dangerously high — only the excess H⁺(aq) present in the stomach is consumed. The neutralisation reaction:

$\text{Mg(OH)}_2\text{(s)} + 2\text{HCl(aq)} \rightarrow \text{MgCl}_2\text{(aq)} + 2\text{H}_2\text{O(l)}$Mg(OH)2​(s)+2HCl(aq)→MgCl2​(aq)+2H2​O(l)

(or in ionic form, $\text{Mg(OH)}_2\text{(s)} + 2\text{H}^+\text{(aq)} \rightarrow \text{Mg}^{2+}\text{(aq)} + 2\text{H}_2\text{O(l)}$Mg(OH)2​(s)+2H+(aq)→Mg2+(aq)+2H2​O(l)).

Used to treat indigestion and heartburn. The same logic explains why Mg(OH)₂ is preferred over NaOH (a soluble strong base) — NaOH would dissolve and raise the gastric pH catastrophically high (chemical burn risk); Mg(OH)₂'s insolubility is a feature, not a bug, providing slow self-limiting neutralisation.

2. Ca(OH)₂ — "Slaked lime" for liming acidic soils

Agricultural soils acidify over time due to (i) acid rain, (ii) nitrification of fertilisers (NH₄⁺ → NO₃⁻ + 2H⁺), and (iii) organic matter decay releasing H⁺. Acidic soils inhibit nutrient uptake by crops and reduce yields. Ca(OH)₂ — known as slaked lime — is spread on fields to neutralise this acidity:

$\text{Ca(OH)}_2\text{(s)} + 2\text{H}^+\text{(aq)} \rightarrow \text{Ca}^{2+}\text{(aq)} + 2\text{H}_2\text{O(l)}$Ca(OH)2​(s)+2H+(aq)→Ca2+(aq)+2H2​O(l)

Ca(OH)₂ is preferred over NaOH (or KOH) for soil liming because it is:

• Cheap — produced by slaking CaO (from heating CaCO₃ / limestone), which is abundant.
• Solid and safer to handle — easier to bag, spread, and store than corrosive liquid bases.
• Only moderately soluble — releases OH⁻ slowly, raising the soil pH gradually and avoiding shock to plants.
• Non-polluting — Ca²⁺ is a beneficial plant nutrient (component of cell walls).

In agricultural contexts, the precursor CaO ("quicklime") is sometimes used directly — it reacts with soil moisture in situ: $\text{CaO(s)} + \text{H}_2\text{O(l)} \rightarrow \text{Ca(OH)}_2\text{(aq)}$CaO(s)+H2​O(l)→Ca(OH)2​(aq) and then the Ca(OH)₂ neutralises soil acidity.

3. CaCO₃ (limestone) — flue-gas desulfurisation (FGD)

Coal-fired power stations and large industrial furnaces burn sulfur-containing fossil fuels, producing SO₂(g) as a major pollutant. Released untreated, SO₂ dissolves in atmospheric moisture to form H₂SO₃ and (after oxidation) H₂SO₄ — the principal cause of acid rain, which damages forests, lakes, buildings, and crops.

Flue-gas desulfurisation (FGD) removes SO₂ by spraying powdered CaCO₃ slurry (or Ca(OH)₂ slurry) into the flue-gas stream. The acidic SO₂ reacts with the alkaline calcium compound to form solid calcium sulfite or sulfate, which is collected:

$\text{CaCO}_3\text{(s)} + \text{SO}_2\text{(g)} \rightarrow \text{CaSO}_3\text{(s)} + \text{CO}_2\text{(g)}$CaCO3​(s)+SO2​(g)→CaSO3​(s)+CO2​(g)

In the presence of excess air, the sulfite is oxidised to sulfate:

$2\text{CaSO}_3\text{(s)} + \text{O}_2\text{(g)} \rightarrow 2\text{CaSO}_4\text{(s)}$2CaSO3​(s)+O2​(g)→2CaSO4​(s)

Or directly with calcium hydroxide slurry:

$\text{Ca(OH)}_2\text{(aq)} + \text{SO}_2\text{(g)} \rightarrow \text{CaSO}_3\text{(s)} + \text{H}_2\text{O(l)}$Ca(OH)2​(aq)+SO2​(g)→CaSO3​(s)+H2​O(l)

Modern coal-fired power stations achieve >95 % SO₂ removal by FGD. The solid CaSO₄ byproduct ("synthetic gypsum") is sold for use in plasterboard and cement manufacture, making FGD essentially cost-neutral for the operator.

4. BaSO₄ — barium meals in medical imaging

BaSO₄ is essentially insoluble in water and in dilute acids (including the HCl of the stomach). Ba²⁺ ions are extremely toxic — they interfere with K⁺ channels in nerve and muscle cells, causing cardiac arrhythmias and paralysis at low doses. But because BaSO₄ does not dissolve, Ba²⁺ ions are never released into the bloodstream.

Barium meals: a patient swallows a thick BaSO₄ suspension before X-ray imaging of the digestive tract. BaSO₄'s high atomic number (Ba: Z = 56) makes it opaque to X-rays, so the suspension coats the inner surface of the stomach and intestines, producing a clear silhouette of the gut on the X-ray image. Any obstructions, ulcers, or anatomical anomalies are immediately visible.

The crucial safety logic:

• Soluble Ba²⁺ salts (BaCl₂, Ba(NO₃)₂) are lethal — they release toxic Ba²⁺ into the bloodstream.
• Insoluble BaSO₄ is safe — Ba²⁺ stays locked in the crystal lattice and passes harmlessly through the gut.

This is one of the cleanest examples in chemistry of how insolubility is engineered for biomedical safety.

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Reaction trees and limewater test for CO₂

flowchart TD
  A[Ca metal] -->|burn in O2| B[CaO 'quicklime']
  B -->|add water 'slaking'| C[Ca OH 2 'slaked lime' / 'limewater']
  C -->|bubble CO2| D[CaCO3 s white precipitate]
  D -->|heat above 825 C| B
  C -->|add dilute HCl| E[CaCl2 aq + H2O l]
  B -->|spray into flue gas| F[CaCO3 / CaSO3 / CaSO4 FGD byproducts]

The CaCO₃ precipitate in limewater is the classic test for CO₂ gas:

$\text{Ca(OH)}_2\text{(aq)} + \text{CO}_2\text{(g)} \rightarrow \text{CaCO}_3\text{(s)} + \text{H}_2\text{O(l)}$Ca(OH)2​(aq)+CO2​(g)→CaCO3​(s)+H2​O(l)

The solution becomes milky / cloudy as fine CaCO₃ particles precipitate. Continued bubbling of CO₂ eventually re-dissolves the CaCO₃ as soluble calcium hydrogencarbonate:

$\text{CaCO}_3\text{(s)} + \text{CO}_2\text{(g)} + \text{H}_2\text{O(l)} \rightarrow \text{Ca(HCO}_3\text{)}_2\text{(aq)}$CaCO3​(s)+CO2​(g)+H2​O(l)→Ca(HCO3​)2​(aq)

The reversal explains both the geological cycle (limestone caves carved by CO₂-rich water) and the requirement to use a brief CO₂ test in school labs (over-bubbling clears the cloudiness and confuses students).

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Worked examples

Example 1 — antacid vs strong base

Q: Explain why milk of magnesia (Mg(OH)₂) is preferred to sodium hydroxide as an antacid.

Answer: NaOH is a strong, highly soluble alkali. Swallowing it would release large amounts of OH⁻(aq), raising the stomach pH from its natural ~1.5 well above 7 — potentially burning gastric tissue and disrupting digestion. Mg(OH)₂ is only sparingly soluble; it remains as a fine suspension that reacts with the excess H⁺(aq) present in the stomach as it dissolves slowly, with the reaction $\text{Mg(OH)}_2\text{(s)} + 2\text{H}^+\text{(aq)} \rightarrow \text{Mg}^{2+}\text{(aq)} + 2\text{H}_2\text{O(l)}$Mg(OH)2​(s)+2H+(aq)→Mg2+(aq)+2H2​O(l). The pH rises modestly (to around 3–4) and stops there, because once the excess H⁺ is consumed there is no driving force to dissolve more Mg(OH)₂. The insolubility provides a self-limiting neutralisation, which is exactly what you want from an antacid.

Example 2 — barium meals

Q: Why is BaSO₄ used in barium meals despite Ba²⁺ being highly toxic?