Why Plaster of Paris Sets Hard Within Minutes of Adding Water

Calcium hardens bones and teeth, magnesium sits at the heart of every chlorophyll molecule, and their compounds build homes, set broken arms and whiten walls before festivals. Group 2 elements resemble Group 1, but with two valence electrons their chemistry follows its own trends.

This lesson covers the general trends of Group 2 elements, calcium oxide, calcium carbonate and Plaster of Paris, and the biological importance of sodium, potassium, magnesium and calcium.

**The Group 2 alkaline earth metals — beryllium, magnesium, calcium, strontium and barium — have an $n{s}^2$ configuration and form ${\mathrm{M}}^{2+}$ ions; they are smaller, harder and less reactive than the alkali metals, and their reactivity increases down the group.

Physical trends:

- First ionisation enthalpy is higher than in Group 1 and decreases down the group
- Flame colours: calcium brick red, strontium crimson and barium apple green; beryllium and magnesium hold their electrons too tightly to colour a flame

Chemical trends:

- Reaction with water: beryllium does not react, magnesium reacts with hot water, and calcium, strontium and barium react with cold water
- Hydroxides become more basic and more soluble down the group
- Sulphates become less** soluble down the group: $\mathrm{BeS}{\mathrm{O}}_4$ and $\mathrm{MgS}{\mathrm{O}}_4$ dissolve readily, but $\mathrm{BaS}{\mathrm{O}}_4$ is almost insoluble
- Carbonates become more stable to heat down the group

Why sulphate solubility falls. The large sulphate ion keeps the lattice enthalpy nearly constant, but hydration enthalpy falls as the cation grows, so less energy is available to break up the lattice.

Anomalous beryllium. Its very small size makes its compounds largely covalent and its oxide amphoteric, and it resembles aluminium through a diagonal relationship.

An everyday example. A barium meal swallowed before a stomach X-ray is safe because barium sulphate is so insoluble, while Epsom salt, magnesium sulphate, dissolves easily in bath water.

The substance. Soluble barium salts are poisonous — barium sulphate is harmless only because almost none of it dissolves.

Calcium oxide is made by heating limestone, calcium carbonate by passing carbon dioxide into lime water, and Plaster of Paris by heating gypsum to 393 K — and all three are workhorses of construction and industry.

Calcium oxide, CaO (quicklime):

- Preparation by heating limestone in a kiln at 1070 to 1270 K, removing carbon dioxide so the reaction goes to completion: $\mathrm{CaC}{\mathrm{O}}_3\rightleftharpoons \mathrm{CaO}+\mathrm{C}{\mathrm{O}}_2$
- Slaking with water gives slaked lime with a hissing sound and much heat: $\mathrm{CaO}+{\mathrm{H}}_2\mathrm{O}\to \mathrm{Ca}(\mathrm{OH}{)}_2$
- A basic oxide that combines with acidic oxides: $\mathrm{CaO}+\mathrm{Si}{\mathrm{O}}_2\to \mathrm{CaSi}{\mathrm{O}}_3$
- Uses: cement, purifying sugar, drying ammonia gas and making bleaching powder

**Calcium carbonate, $\mathrm{CaC}{\mathrm{O}}_3$ (limestone, marble, chalk):

- Preparation**: $\mathrm{Ca}(\mathrm{OH}{)}_2+\mathrm{C}{\mathrm{O}}_2\to \mathrm{CaC}{\mathrm{O}}_3+{\mathrm{H}}_2\mathrm{O}$; excess carbon dioxide redissolves it as soluble calcium hydrogencarbonate
- Properties: almost insoluble in water; releases carbon dioxide with acids
- Uses: building stone, flux in metal extraction, antacid tablets and toothpaste filler

**Plaster of Paris, $\mathrm{CaS}{\mathrm{O}}_4\cdot \frac{1}{2}{\mathrm{H}}_2\mathrm{O}$:

- Preparation** by heating gypsum at 393 K:

$$2(\mathrm{CaS}{\mathrm{O}}_4\cdot 2{\mathrm{H}}_2\mathrm{O})\to 2(\mathrm{CaS}{\mathrm{O}}_4)\cdot {\mathrm{H}}_2\mathrm{O}+3{\mathrm{H}}_2\mathrm{O}$$

- Setting: mixed with water, it turns back into interlocking gypsum crystals and hardens within minutes, expanding slightly
- Heated well above 393 K, it becomes anhydrous dead burnt plaster, which no longer sets
- Uses: casts for fractures, dental moulds, statues and false ceilings

Worked example. From $\mathrm{CaC}{\mathrm{O}}_3\to \mathrm{CaO}+\mathrm{C}{\mathrm{O}}_2$, 100 g of limestone gives 56 g of quicklime and 44 g of carbon dioxide.

An everyday example. Whitewashing walls with lime before Diwali leaves a bright finish because slaked lime slowly absorbs carbon dioxide from the air and forms a thin, hard layer of calcium carbonate.

The substance. Plaster of Paris sets by rehydration, not by drying — it needs water to rebuild gypsum crystals, which is why overheated plaster that has lost all its water is useless.

Sodium and potassium ions carry nerve signals and control the balance of water and charge across cell membranes, magnesium ions activate enzymes and sit at the centre of chlorophyll, and calcium builds bones and teeth and triggers muscle contraction and blood clotting.

Sodium and potassium:

- Sodium ions are found mainly outside cells, in blood plasma and the fluid around cells; potassium ions are found mainly inside cells
- The sodium-potassium pump uses energy from ATP to move sodium ions out of cells and potassium ions in
- The resulting difference in ion concentrations lets nerve impulses travel and controls the movement of water into and out of cells

Magnesium:

- The central atom of chlorophyll, so plants cannot photosynthesise without it
- Needed by enzymes that use ATP, and so for releasing energy from food

Calcium:

- Present in bones and teeth as calcium phosphate, mainly hydroxyapatite, $\mathrm{C}{\mathrm{a}}_{10}(\mathrm{P}{\mathrm{O}}_4{)}_6(\mathrm{OH}{)}_2$
- Calcium ions trigger muscle contraction, including the heartbeat, and are essential for blood clotting

An everyday example. Oral rehydration solution given to children with diarrhoea contains sodium and potassium salts with glucose, replacing the ions lost from the body so that cells can hold water again.

The substance. Sodium and potassium are not interchangeable in the body — though chemically similar, one belongs outside cells and the other inside, and that difference is what powers nerve signals.