The Elements
Lithium, sodium, potassium, rubidium, caesium and francium make up Group I — the alkali metals, named for the alkaline solutions they form in water. Each has a single valence electron and a constant +1 oxidation state, which it loses readily to form ionic bonds — they’re highly electropositive, and that electropositivity (and so their reactivity) increases down the group. They’re silvery-white, soft enough to cut with a knife, and good conductors of heat and electricity — but so reactive that they’re stored under paraffin oil and are always found combined in nature, extracted industrially by electrolysing their molten chlorides.
Chemical Properties
- Tarnishing in air: exposed to air, they tarnish through a succession of oxide, hydroxide and carbonate. 4Na(s) + O₂(g) → 2Na₂O(s); Na₂O(s) + H₂O(l) → 2NaOH(aq); 2NaOH(aq) + CO₂(g) → Na₂CO₃(aq) + H₂O(l)
- Burning in oxygen: produces characteristic flame colours — sodium burns with a bright golden-yellow flame, potassium with a lilac one — each forming a white solid that dissolves in water to an alkaline solution. 4K(s) + O₂(g) → 2K₂O(s); K₂O(s) + H₂O(l) → 2KOH(aq)
- Reaction with cold water: vigorous, evolving hydrogen and forming the alkaline hydroxide solution that gives the group its name. 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g); 2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g)
- Reaction with dilute acids: explosively vigorous (too dangerous to demonstrate directly), liberating hydrogen. 2Na(s) + 2HCl(aq) → 2NaCl(aq) + H₂(g)
- Reaction with hydrogen: forms ionic (electrovalent) hydrides. 2Na(s) + H₂(g) → 2NaH(s); 2K(s) + H₂(g) → 2KH(s)
- Effect of heat on nitrates: decompose to the nitrite plus oxygen. 2NaNO₃(s) → 2NaNO₂(s) + O₂(g); 2KNO₃(s) → 2KNO₂(s) + O₂(g)
- Effect of heat on carbonates and hydroxides: Group I carbonates and hydroxides are thermally stable — Na₂CO₃, K₂CO₃, NaOH and KOH all resist decomposition on heating.
Lithium’s Anomalous Behaviour
Lithium is the exception to several of the patterns above — its small, strongly polarising ion makes it behave more like Group II’s magnesium and calcium (a pattern chemists call a “diagonal relationship”). It’s harder and less reactive than the rest of the group, tarnishing and reacting with water only slowly. Its chloride, LiCl, is deliquescent, like CaCl₂. Unlike the rest of Group I, its carbonate and nitrate both decompose on heating instead of resisting it — matching the Group II pattern rather than its own group’s: Li₂CO₃(s) → Li₂O(s) + CO₂(g); 4LiNO₃(s) → 2Li₂O(s) + 4NO₂(g) + O₂(g). Its hydroxide, unlike NaOH and KOH, is only moderately soluble and not strongly caustic. Its hydrogencarbonate, LiHCO₃, exists only in solution, as with the Group II hydrogencarbonates. And like magnesium and calcium, it reacts directly with nitrogen to form a nitride: 6Li(s) + N₂(g) → 2Li₃N(s).
Sodium: Extraction and Properties
Sodium occurs naturally as sodium chloride — in rock salt deposits and dissolved in seawater — and is too reactive to occur free. It’s extracted by electrolysing molten sodium chloride in a Down’s cell (a fire-brick-lined steel container); since pure NaCl melts at a inconveniently high 801°C, calcium chloride is mixed in to lower the melting point to a more workable 600°C. The anode is graphite, the cathode steel, and a steel gauze diaphragm keeps the chlorine gas produced at the anode away from the sodium forming at the cathode.
Cathode: Na⁺(l) + e⁻ → Na(s). Anode: 2Cl⁻(l) → Cl₂(g) + 2e⁻
Physical properties: a silvery-white metal, soft enough to cut with a knife, low melting point (~97°C), a good conductor, floats on water, burns with a golden-yellow flame.
Chemical properties: burns in excess oxygen to sodium peroxide (4Na(s) + 2O₂(g) → 2Na₂O₂(s)), or in limited air to the ordinary oxide (4Na(s) + O₂(g) → 2Na₂O(s)); combines with most non-metals on heating (2Na(s) + H₂(g) → 2NaH(s); 2Na(s) + S(s) → Na₂S(s)); reacts violently with cold water; amalgamates with mercury (2Na(s) + Hg(l) → 2Na/Hg(l)).
Sodium Hydroxide
Manufacture: electrolysis of brine using a mercury-cathode (Kellner-Solvay) cell — see the Electrolysis post for the full mechanism. Physical properties: melts at 320°C; a deliquescent, crystalline solid; dissolves in water with a lot of heat released.
Chemical properties:
- Absorbs CO₂ from the air to form sodium carbonate — which is why it should never be left uncorked. 2NaOH(aq) + CO₂(g) → Na₂CO₃(aq) + H₂O(l)
- Reacts with acidic oxides to give a salt and water. 2NaOH(aq) + SO₂(g) → Na₂SO₃(aq) + H₂O(l)
- Releases ammonia when heated with an ammonium salt. NaOH(aq) + NH₄Cl(s) → NaCl(aq) + H₂O(l) + NH₃(g)
- Attacks glass, forming sodium silicate, which reacts further with CO₂ to form sodium carbonate decahydrate — this can cement a glass stopper firmly into a bottle neck (removable with HCl).
- Reacts with amphoteric metals like aluminium and zinc, releasing hydrogen. 2Al(s) + 2NaOH(aq) + 6H₂O(l) → 2NaAl(OH)₄(aq) + 3H₂(g)
- Precipitates insoluble hydroxides from many metal salt solutions. 2NaOH(aq) + FeCl₂(aq) → Fe(OH)₂(s) + 2NaCl(aq); 2NaOH(aq) + CuSO₄(aq) → Cu(OH)₂(s) + Na₂SO₄(aq)
Uses: purifying bauxite in aluminium extraction; soap manufacture; paper and fibre industries; baking powder.
Sodium Carbonate
Prepared by bubbling excess CO₂ into NaOH solution to form sodium hydrogencarbonate, filtering off the solid, then heating it: NaOH(aq) + CO₂(g) → NaHCO₃(aq); 2NaHCO₃(s) → Na₂CO₃(s) + H₂O(l) + CO₂(g). Properties: a soluble, hygroscopic white powder (“soda ash”); the hydrated form, washing soda (Na₂CO₃·10H₂O), is a translucent crystalline solid that’s efflorescent (loses water of crystallisation in air, becoming a monohydrate); reacts with acids to release CO₂. Uses: seasoning; softening hard water (washing soda); baking powder (sodium hydrogencarbonate).
