Aluminium and d-block metals (Fe, Cu, Zn, Cr, Mn)

d-block elements (transition metals) are the elements of the B groups of the periodic table, whose atoms fill the d-orbitals of the shell below the outer one. The most important in practice are iron, copper, zinc, chromium and manganese; aluminium is studied alongside them because, although it is a p-block element, its amphoteric behaviour and industrial importance make it similar. These metals share variable oxidation states, coloured compounds, amphoteric oxides and hydroxides, and the ability to form complex compounds. Exactly these properties lead to exam questions on colour changes, on "which salt forms", and on mixtures.

§1. General features of the d-block

With few outer electrons, all d-block elements are metals. Their valence electrons include the d-electrons as well as the outer s-electrons, so they show many, variable oxidation states: manganese from +2 to +7, chromium +2, +3, +6, iron +2, +3. A partly filled d subshell makes compounds coloured: CuSO₄·5H₂O is blue, FeCl₃ yellow-brown, KMnO₄ purple, K₂Cr₂O₇ orange. Ions with a full d¹⁰ subshell (Zn²⁺) are colourless.

A general rule: as an element’s oxidation state rises, its oxides and hydroxides change from basic to amphoteric and then to acidic; compounds in low oxidation states are reducing agents, those in high states oxidising agents.

§2. Aluminium

Aluminium is a reactive metal, but in air it is covered by a thin, dense film of Al₂O₃, so it does not react with water and resists corrosion. Cold conc. HNO₃ and H₂SO₄ passivate it.
Amphoteric behaviour — aluminium, its oxide and its hydroxide react with both acids and alkalis:

When alkali is added drop by drop to an aluminium salt solution, a white gelatinous precipitate of Al(OH)₃ forms first (3 mol OH⁻ per mol Al³⁺); with excess alkali (the fourth mole) it dissolves.
The thermite reaction: Fe₂O₃ + 2Al → Al₂O₃ + 2Fe releases a great deal of heat and is used to weld rails. Aluminium is produced industrially by electrolysis of Al₂O₃ dissolved in molten cryolite.

Worked example

5.4 g of aluminium dissolved in excess sodium hydroxide solution. What volume of hydrogen (STP) is released?

1. n(Al) = 5.4/27 = 0.2 mol.
2. 2Al → 3H₂: n(H₂) = 0.2·3/2 = 0.3 mol.
3. V = 0.3·22.4 = 6.72 L.

§3. Iron

Iron is a moderately reactive metal and the second most abundant metal in the Earth’s crust. Depending on the strength of the oxidising agent it forms two ions:

• with HCl, dilute H₂SO₄ or copper salts — Fe²⁺: Fe + 2HCl → FeCl₂ + H₂; Fe + CuSO₄ → FeSO₄ + Cu.
• with Cl₂, hot concentrated acids or HNO₃ — Fe³⁺: 2Fe + 3Cl₂ → 2FeCl₃.
• burning in oxygen gives the mixed oxide Fe₃O₄ (FeO·Fe₂O₃): 3Fe + 2O₂ → Fe₃O₄.

Fe(OH)₂ is a greenish-white precipitate that turns brown in air: 4Fe(OH)₂ + O₂ + 2H₂O → 4Fe(OH)₃. Fe(OH)₃ is a brown precipitate, weakly amphoteric.

Alloys: cast iron (more than 2 % carbon) and steel (less than 2 %). In the blast furnace iron ore is reduced by carbon monoxide: Fe₂O₃ + 3CO → 2Fe + 3CO₂. Rusting is electrochemical corrosion involving moisture and oxygen.

Worked example

11.2 g of iron reacted: a) with excess hydrochloric acid; b) with excess chlorine. Find the mass of salt in each case.

1. n(Fe) = 11.2/56 = 0.2 mol.
2. a) FeCl₂ forms: m = 0.2·127 = 25.4 g.
3. b) FeCl₃ forms: m = 0.2·162.5 = 32.5 g.

§4. Copper and zinc

Copper ([Ar] 3d¹⁰ 4s¹, +1 and +2) lies below hydrogen in the reactivity series, so it does not react with HCl or dilute H₂SO₄; it dissolves only in oxidising acids: Cu + 2H₂SO₄(conc.) → CuSO₄ + SO₂ + 2H₂O. Compounds: CuO (black), Cu₂O (red), Cu(OH)₂ (pale blue precipitate that decomposes on heating to CuO and water), CuSO₄·5H₂O (blue vitriol; the anhydrous salt is white). In air copper objects acquire a green coating of basic copper carbonate (CuOH)₂CO₃. Alloys: brass (Cu + Zn) and bronze (Cu + Sn).

Zinc ([Ar] 3d¹⁰ 4s², only +2) is an amphoteric metal: Zn + 2HCl → ZnCl₂ + H₂; Zn + 2NaOH + 2H₂O → Na₂[Zn(OH)₄] + H₂. ZnO and Zn(OH)₂ are amphoteric too. Coating iron with zinc (galvanising) is widely used against corrosion: zinc is more reactive, so it corrodes instead of the iron.

§5. Chromium and manganese

For these metals, properties change from basic to amphoteric to acidic as the oxidation state rises:

Chromium is the hardest metal and makes steel stainless (stainless steel contains chromium and nickel). The chromate–dichromate equilibrium depends on the medium:

Dichromate is stable in acid, chromate in alkali. K₂Cr₂O₇ and KMnO₄ are strong oxidising agents; the reduction product of KMnO₄ depends on the medium: Mn²⁺ (colourless) in acid, MnO₂ (brown precipitate) in neutral solution and MnO₄²⁻ (green) in alkali. Laboratory oxygen: 2KMnO₄ → K₂MnO₄ + MnO₂ + O₂; laboratory chlorine: MnO₂ + 4HCl → MnCl₂ + Cl₂ + 2H₂O.

§6. An official test item

97.6 g of a mixture of cadmium nitrate and iron was heated in air and its mass did not change (CdO and Fe₂O₃ form). Find the mass of the salt in the mixture.

Worked example

1. 2Cd(NO₃)₂ → 2CdO + 4NO₂ + O₂: 1 mol of salt (236 g) leaves 128 g of CdO, a loss of 108 g.
2. 4Fe + 3O₂ → 2Fe₂O₃: 1 mol of Fe (56 g) gives 80 g of Fe₂O₃, a gain of 24 g.
3. The mass is unchanged: 108x = 24y → y = 4.5x; 236x + 56·4.5x = 488x = 97.6 → x = 0.2.
4. Mass of salt: 0.2·236 = 47.2 g.

Check yourself

1. 2.7 g of Al dissolved in excess hydrochloric acid. What volume of H₂ (STP) is released?

2. Which reagent converts iron into FeCl₃?

3. Which reagent detects the Fe³⁺ ion?

4. What kind of oxide is CrO₃?

5. 0.4 mol of NaOH was added to 0.1 mol of AlCl₃. Does a precipitate remain?

6. What forms when Cu(OH)₂ is heated?

7. What is the mass fraction of iron in Fe₃O₄?

8. What is the oxidation state of manganese in KMnO₄?

Common mistakes

• Writing that iron gives FeCl₃ with hydrochloric acid: FeCl₂ forms.
• Forgetting that Al(OH)₃ dissolves in excess alkali.
• Thinking aluminium reacts readily with water: its oxide film protects it.
• Treating Fe₃O₄ as an oxide of Fe³⁺ only: it contains both Fe²⁺ and Fe³⁺.
• Writing that copper displaces hydrogen from dilute H₂SO₄.