The Elements
Beryllium, magnesium, calcium, strontium, barium and radium make up Group II. Each has two valence electrons and a constant +2 oxidation state, lost to form ionic bonds — they’re reactive, though less so than Group I, and their reactivity likewise increases down the group. Their ions are what make water hard.
Chemical Properties
- Reaction with oxygen: burn to their oxides. 2Mg(s) + O₂(g) → 2MgO(s); 2Ca(s) + O₂(g) → 2CaO(s)
- Reaction with water: calcium reacts with cold water to a weak alkaline solution, evolving hydrogen: Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g). Magnesium is unreactive with cold water but reacts with steam, giving the insoluble oxide instead of the hydroxide: Mg(s) + H₂O(l) → MgO(s) + H₂(g).
- Effect of heat on nitrates: decompose to the oxide, nitrogen dioxide and oxygen. 2Mg(NO₃)₂(s) → 2MgO(s) + 4NO₂(g) + O₂(g); 2Ca(NO₃)₂(s) → 2CaO(s) + 4NO₂(g) + O₂(g)
- Effect of heat on hydroxides: decompose to the oxide and water. Ca(OH)₂(s) → CaO(s) + H₂O(g)
- Effect of heat on carbonates: all Group II carbonates decompose to the oxide and CO₂ (they’re insoluble to begin with). CaCO₃(s) → CaO(s) + CO₂(g)
- Reaction with nitrogen: forms nitrides on heating. 3Mg(s) + N₂(g) → Mg₃N₂(s); 3Ca(s) + N₂(g) → Ca₃N₂(s). Warmed with water, the nitride hydrolyses to the hydroxide and ammonia: Mg₃N₂(s) + 6H₂O(l) → 3Mg(OH)₂(s) + 2NH₃(g)
- Reaction with hydrogen: forms ionic hydrides. Ca(s) + H₂(g) → CaH₂(s); Mg(s) + H₂(g) → MgH₂(s)
Calcium and Its Compounds
Calcium occurs abundantly as calcium carbonate — limestone, marble, chalk, calcite — and as calcium phosphate in bone and teeth. It’s a silvery-white metal, denser than water, melting at 851°C, and burns in air with a brick-red flame to a mixture of oxide and nitride: 2Ca(s) + O₂(g) → 2CaO(s); 3Ca(s) + N₂(g) → Ca₃N₂(s). It reacts with dilute acids, displacing hydrogen: Ca(s) + 2HCl(aq) → CaCl₂(aq) + H₂(g).
Calcium carbonate: an insoluble white solid (except in CO₂-containing water); releases CO₂ with dilute acid (CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)); dissolves in CO₂-laden water to the soluble bicarbonate responsible for temporary hardness (CaCO₃(s) + H₂O(l) + CO₂(g) → Ca(HCO₃)₂(aq)); decomposes to the oxide above about 1000°C. Uses: iron extraction; cement and glass manufacture; commercial quicklime; road-building; whitening paper; neutralising acidic soils and acid-rain-affected lakes.
Calcium oxide (quicklime): a hygroscopic white solid, melting point 2600°C, glowing brilliantly at very high temperatures (“limelight”). A strong base — reacts vigorously with acids (CaO(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l)) and, famously, with water: swelling, crackling and hissing as it releases heat to form slaked lime (CaO(s) + H₂O(l) → Ca(OH)₂(s) + heat — “the slaking of lime”). Releases ammonia from ammonium salts: CaO(s) + 2NH₄Cl(s) → CaCl₂(aq) + H₂O(l) + 2NH₃(g). Uses: neutralising acids; making slaked lime; drying agent for ammonia.
Calcium hydroxide (slaked lime as a solid; limewater in solution), CaO(s) + H₂O(l) → Ca(OH)₂(aq): a slightly soluble white solid, weakly alkaline in solution. Decomposes on heating back to the oxide, and famously turns milky with CO₂ (Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l)) — clearing again with excess CO₂ as the soluble bicarbonate reforms. Uses: testing for CO₂ and for water; making whitewash.
Hardness of Water — Supplement
Rain dissolves atmospheric CO₂ into a weakly acidic solution, which then dissolves limestone as it percolates through rock, picking up Ca(HCO₃)₂ — this is the main natural origin of hard water. Some rocks also contain slightly soluble minerals like gypsum (CaSO₄·2H₂O) and kieserite (MgSO₄·2H₂O), contributing to permanent hardness.
Additional softening methods, beyond boiling and adding calcium hydroxide or washing soda (Part 2): ion exchange passes water through a resin whose bound sodium ions swap for the dissolved calcium (2Na⁺(resin) + Ca²⁺(aq) → Ca²⁺(resin) + 2Na⁺(aq), a 2:1 exchange required by charge balance); distillation removes all dissolved ions but is expensive at scale; and modern detergent softeners contain phosphates that tie up calcium as insoluble calcium phosphate, Ca₃(PO₄)₂.
Soap and scum: soap is a sodium (or potassium) salt of a long-chain organic acid — sold in Cameroon as “savon azur,” chemically sodium stearate (sodium octadecanoate). It’s soluble, but the calcium salt isn’t: mixed with hard water, the calcium ions react with the soap to form an insoluble precipitate — scum — instead of a lather. Ca²⁺(aq) + 2C₁₇H₃₅COO⁻(aq) → Ca(C₁₇H₃₅COO)₂(s)
A fuller balance of hard and soft water: hard water tastes more pleasant, supplies calcium for bones and teeth, suits beer brewing, and its scale can actually coat lead pipes with a protective layer that reduces lead poisoning risk — but it wastes soap, furs up kettles and pipes, and isn’t ideal for precise laboratory work. Soft water lathers efficiently and suits laboratory use, but is less pleasant to drink, since it lacks those same dissolved minerals.
