Position and General Properties
Transition metals occupy the block between Group II and Group III, from period 4 onward (scandium through zinc form the first, most commonly studied row). They’re sometimes called “heavy metals,” since their densities are much higher than the alkali and alkaline earth metals’.
Physical properties: hard and strong (mercury is the one liquid-at-room-temperature exception); high melting and boiling points; high tensile strength; good conductors of heat and electricity; ductile (drawn into wires) and malleable (beaten into sheets).
Characteristic chemical properties:
- Variable oxidation states — e.g. iron as Fe²⁺ or Fe³⁺ — arising from their incompletely filled d-orbitals.
- Coloured compounds — e.g. CuSO₄ is blue, FeSO₄ is green — especially in hydrated form.
- Formation of complex ions, where a central metal ion accepts electron pairs from surrounding “ligand” atoms or molecules — e.g. [Co(H₂O)₆]²⁺ (hexaaquacobalt(II) ion), [Fe(CN)₆]³⁻ (hexacyanoferrate(III) ion), [CoCl₄]²⁻ (tetrachlorocobaltate(II) ion). (A few main-group elements form complex ions too, e.g. [Al(OH)₄]⁻.)
- Catalytic activity — transition metals and their compounds catalyse many industrial gas reactions: manganese(IV) oxide catalyses hydrogen peroxide’s decomposition (2H₂O₂(aq) → 2H₂O(l) + O₂(g)) and potassium chlorate’s thermal decomposition (2KClO₃(s) → 2KCl(s) + 3O₂(g)); finely divided iron catalyses the Haber process (N₂(g) + 3H₂(g) → 2NH₃(g)); vanadium(V) oxide catalyses the Contact process’s key step (2SO₂(g) + O₂(g) → 2SO₃(g)).
Iron
- Rusting: in moist air, iron gradually forms hydrated iron(III) oxide (rust) — porous, unlike the protective oxide layers on zinc or aluminium, so rusting continues until all the iron is consumed. 4Fe(s) + 3O₂(g) + 2xH₂O(l) → 2Fe₂O₃·xH₂O(s)
- Reaction with steam: red-hot iron decomposes steam, releasing hydrogen. 3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)
- Burning in air: finely divided iron (filings) heated in air forms the magnetic mixed oxide. 3Fe(s) + 2O₂(g) → Fe₃O₄(s)
- Reaction with chlorine: heated iron and dry chlorine gas form black, deliquescent iron(III) chloride crystals (stored in a desiccator). 2Fe(s) + 3Cl₂(g) → 2FeCl₃(s). Iron also combines directly with sulphur, phosphorus and carbon, but not nitrogen. Fe(s) + S(s) → FeS(s)
- Reaction with dilute acids: dissolves readily, giving iron(II) salts and hydrogen. Fe(s) + 2HCl(aq) → FeCl₂(aq) + H₂(g); Fe(s) + H₂SO₄(aq) → FeSO₄(aq) + H₂(g)
- Reaction with dilute nitric acid: gives a mix of products depending on conditions — iron(II) nitrate, dinitrogen oxide and water is one balanced outcome: 4Fe(s) + 10HNO₃(aq) → 4Fe(NO₃)₂(aq) + N₂O(g) + 5H₂O(l). Concentrated nitric acid, by contrast, doesn’t react at all — it oxidises the surface to a protective layer of oxide, rendering the iron “passive,” which is why concentrated nitric acid is safely stored in iron containers.
Testing for iron ions with sodium hydroxide (aqueous ammonia gives similar results): Fe²⁺ gives a dirty green gelatinous precipitate, insoluble in excess — Fe²⁺(aq) + 2NaOH(aq) → Fe(OH)₂(s) + 2Na⁺(aq). Fe³⁺ gives a reddish-brown gelatinous precipitate, also insoluble in excess — Fe³⁺(aq) + 3NaOH(aq) → Fe(OH)₃(s) + 3Na⁺(aq).
Copper
- Reaction with air: copper sits very low on the reactivity series and is stable in pure, dry air. In moist air it slowly develops a green coating of basic copper carbonate (and, in moist sea air, basic copper chloride). Heated in air, it oxidises to black copper(II) oxide: 2Cu(s) + O₂(g) → 2CuO(s) — or, in some conditions, red-to-black copper(I) oxide: 4Cu(s) + O₂(g) → 2Cu₂O(s).
- Reaction with acids: too low in the reactivity series to displace hydrogen from ordinary dilute acids, but attacked by oxidising acids. With concentrated sulphuric acid: Cu(s) + 2H₂SO₄(conc)(aq) → CuSO₄(aq) + 2H₂O(l) + SO₂(g). With nitric acid, the product depends on concentration — dilute acid gives nitrogen monoxide, concentrated gives nitrogen dioxide: 3Cu(s) + 8HNO₃(aq) → 3Cu(NO₃)₂(aq) + 4H₂O(l) + 2NO(g); Cu(s) + 4HNO₃(aq) → Cu(NO₃)₂(aq) + 2H₂O(l) + 2NO₂(g).
- Displacement: being near the bottom of the reactivity series, copper is displaced from solutions of its salts by most other metals.
Testing for Cu²⁺ with sodium hydroxide: a blue gelatinous precipitate, insoluble in excess — Cu²⁺(aq) + 2NaOH(aq) → Cu(OH)₂(s) + 2Na⁺(aq). All hydrated copper salts are blue or green, dissolving to blue solutions.
Copper(I) vs copper(II): copper forms both, by losing one or two electrons respectively, but +2 is far more common — copper(I) is unstable in aqueous solution, spontaneously converting to a mix of copper(II) and copper metal (disproportionation, meaning simultaneous self-oxidation and self-reduction): 2Cu⁺(aq) → Cu²⁺(aq) + Cu(s). Copper(I) chloride is a white covalent solid that oxidises to green copper(II) chloride in moist air; it can be made by passing sulphur dioxide (a reducing agent) through a mixed copper(II) sulphate/sodium chloride solution: 2Cu²⁺(aq) + 2Cl⁻(aq) + SO₂(g) + 2H₂O(l) → 2CuCl(s) + SO₄²⁻(aq) + 4H⁺(aq). Copper(I) oxide forms as a red precipitate when an alkaline copper(II) solution is reduced — the basis of the Fehling’s and Benedict’s tests for reducing sugars.
