Chemistry · Chemistry of the elements · Grade 9

Iron and its compounds

How do we tell Fe²⁺ from Fe³⁺, and what does the iron in our blood do?

The situation

The cast-iron cauldron in the kitchen, the steel beams of bridges and buildings — all are made of iron. When iron rusts it turns brown, and its salts come in two kinds: a Fe²⁺ solution is pale green, a Fe³⁺ solution is yellow-brown. Even the red of blood is linked with iron — the iron in haemoglobin picks up oxygen in the lungs and carries it to the tissues. How do we tell the two kinds of iron ion apart, and how much iron is there in our blood?

Topic

Iron and its compounds. Iron: Fe²⁺ solution is pale green, Fe³⁺ yellow-brown; a dark blue precipitate with red prussiate and the iron in blood

Key ideas

Iron(II) and iron(III) hydroxides
Qualitative tests
Haemoglobin

Formulas

4 Fe (OH)_{2} + O_{2} + 2 H_{2} O \overset{}{\to} 4 Fe (OH)_{3}

Fe(OH)₂ is oxidised in air: the white-green precipitate turns into brown Fe(OH)₃ (lab work 15).

Fe^{2 +} + K^{+} + [Fe (CN)_{6}]^{3 -} \overset{}{\to} KFe [Fe (CN)_{6}]

Fe²⁺ + red prussiate → dark blue precipitate (Turnbull's blue).

Fe^{3 +} + K^{+} + [Fe (CN)_{6}]^{4 -} \overset{}{\to} KFe [Fe (CN)_{6}]

Fe³⁺ + yellow prussiate → dark blue precipitate (Prussian blue) — in fact the same substance.

Fe^{3 +} + SCN^{-} \overset{}{⇌} [Fe (SCN)]^{2 +}

Fe³⁺ + thiocyanate → blood-red solution (lab work 16).

m (Fe) = m (Hb) \cdot \frac{4 A_{r} (Fe)}{M (Hb)}

Iron in haemoglobin: 4 · 55.85 / 64 500 ≈ 0.35 %.

Theory

Iron: position and properties (§34)

Iron, Fe, is element 26, in period 4, group VIIIB (group 8): [Ar] 3d⁶ 4s². In compounds it is mainly +2 and +3 (+6 in ferrates). By mass it makes up ≈ 5 % of the Earth's crust — 4th after oxygen, silicon and aluminium; in nature it occurs mainly in compounds: magnetite Fe₃O₄ (ω(Fe) = 72.4 %), haematite Fe₂O₃ (69.9 %), limonite 2Fe₂O₃·3H₂O (59.8 %), siderite FeCO₃ (48.2 %); sulfuric acid is made from pyrite FeS₂. Pure iron is silvery white and relatively soft, ρ = 7.87 g/cm³, and melts at 1538 °C; up to 770 °C it is ferromagnetic — attracted by a magnet. In industry, alloys of iron are used: cast iron (more than 2.14 % C) and steel (up to 2.14 % C) — see the “Making cast iron and steel” experiment for details.

Chemical properties

Iron is a metal of medium reactivity. It displaces hydrogen from hydrochloric and dilute sulfuric acids, forming iron(II) salts: $Fe + 2 HCl \overset{}{\to} FeCl_{2} + H_{2}$ ; with chlorine, a strong oxidising agent, it forms iron(III) chloride: $2 Fe + 3 Cl_{2} \overset{}{\to} 2 FeCl_{3}$ . It burns in oxygen to give iron scale Fe₃O₄ (FeO·Fe₂O₃): $3 Fe + 2 O_{2} \overset{}{\to} Fe_{3} O_{4}$ ; with sulfur it gives FeS, and with hot steam Fe₃O₄ and H₂: $3 Fe + 4 H_{2} O \overset{}{\to} Fe_{3} O_{4} + 4 H_{2}$ . Cold concentrated nitric and sulfuric acids passivate iron. Iron displaces less reactive metals from their salts: $Fe + CuSO_{4} \overset{}{\to} FeSO_{4} + Cu$ . In damp air it rusts — both oxygen and water are needed: $4 Fe + 3 O_{2} + 6 H_{2} O \overset{}{\to} 4 Fe (OH)_{3}$ (rust is hydrated iron(III) oxide).

Iron hydroxides (lab work 15)

When alkali is added to a solution of an iron(II) salt, Fe(OH)₂ is precipitated: $Fe^{2 +} + 2 OH^{-} \overset{}{\to} Fe (OH)_{2}$ . Pure Fe(OH)₂ is white, but it begins to be oxidised at once by atmospheric oxygen: first white-green, then dark green-grey, and finally brown Fe(OH)₃: $4 Fe (OH)_{2} + O_{2} + 2 H_{2} O \overset{}{\to} 4 Fe (OH)_{3}$ . To slow the oxidation, a layer of oil is poured over the solution. From an iron(III) salt, red-brown Fe(OH)₃ is precipitated at once: $Fe^{3 +} + 3 OH^{-} \overset{}{\to} Fe (OH)_{3}$ . Both hydroxides dissolve in acid ( $Fe (OH)_{3} + 3 HCl \overset{}{\to} FeCl_{3} + 3 H_{2} O$ ); Fe(OH)₃ is almost insoluble in alkali — its amphoteric character is very weak. On heating it decomposes to Fe₂O₃: $2 Fe (OH)_{3} \overset{}{\to} Fe_{2} O_{3} + 3 H_{2} O$ . Fe²⁺ is a reducing agent (in air and with chlorine it passes to Fe³⁺: $2 FeCl_{2} + Cl_{2} \overset{}{\to} 2 FeCl_{3}$ ), and Fe³⁺ is an oxidising agent.

Detecting Fe²⁺ and Fe³⁺ (lab work 16)

Qualitative reactions: Fe²⁺ + red prussiate K₃[Fe(CN)₆] → dark blue precipitate (Turnbull's blue): $Fe^{2 +} + K^{+} + [Fe (CN)_{6}]^{3 -} \overset{}{\to} KFe [Fe (CN)_{6}]$ . Fe³⁺ + yellow prussiate K₄[Fe(CN)₆] → dark blue precipitate (Prussian blue): $Fe^{3 +} + K^{+} + [Fe (CN)_{6}]^{4 -} \overset{}{\to} KFe [Fe (CN)_{6}]$ . Modern research shows that these two precipitates are the same substance — KFe[Fe(CN)₆]. The “reverse” pairs give no blue precipitate: Fe³⁺ + K₃[Fe(CN)₆] — a brown-green solution, Fe²⁺ + K₄[Fe(CN)₆] — a white precipitate (slowly turning blue in air). Fe³⁺ + potassium thiocyanate KSCN → blood-red solution: $Fe^{3 +} + SCN^{-} \overset{}{⇌} [Fe (SCN)]^{2 +}$ ; with Fe²⁺ there is almost no colour (traces of Fe³⁺ oxidised in air may give a pale pink tint). With alkali: Fe²⁺ — a white-green precipitate, Fe³⁺ — a red-brown precipitate. To be reliable, test an unknown solution with at least two reagents.

Uses and biological role (§35)

Iron is the most widely used metal: steel and cast iron are used for construction, engineering, railways and cookware (the cast-iron cauldron); in Uzbekistan steel is made at the metallurgical works in Bekabad. Its compounds: iron(II) sulfate FeSO₄·7H₂O — for plant protection and as a fertiliser; FeCl₃ — a coagulant in water treatment; Fe₂O₃ — a red pigment; Prussian blue — a blue dye. Biological role: the haemoglobin molecule in red blood cells has four haem groups, and the Fe²⁺ ion at the centre of each binds one O₂ molecule and carries it from the lungs to the tissues. About 0.35 % of the mass of haemoglobin is iron ( $4 \cdot 55.85 / 64 500$ ). An adult has a few grams of iron, of which ≈ 2.5 g is in haemoglobin; ≈ 1 mg a day (1.5–2 mg in women) is lost and replaced by food (meat, liver, pulses, greens) — if there is too little, anaemia develops.

Safety and the model

Safety: work with alkalis and acids wearing safety goggles; if alkali gets in the eye, wash it with plenty of water. The prussiates (hexacyanoferrates) are stable, but do not heat them with acid — poisonous HCN may be released. Thiocyanates are poisonous too — do not taste them. Model: the amount of precipitate follows stoichiometry (the final state depends on the moles added); the oxidation of Fe(OH)₂ in air is simplified (1 s ≈ 1 min, and under the oil layer it is taken not to oxidise). The FeSO₄ solution is taken to contain 1 % Fe³⁺ (the part oxidised in air). Colours are approximate by the Beer–Lambert law. The haem molecule is schematic: the atoms are reduced in size and the protein part of haemoglobin is not shown; M(Hb) ≈ 64 500 g/mol.

Did you know?

Turnbull's blue and Prussian blue were long considered different substances; modern research has shown that both are the same substance — iron(III) hexacyanoferrate(II). Prussian blue was obtained by accident in Berlin at the beginning of the 18th century and was one of the first synthetic pigments.
An adult's body contains a few grams of iron, of which about 2.5 g is in haemoglobin: the four iron ions in each haemoglobin molecule carry four O₂ molecules.

Inside the experiment

The experiment starts from a real-life situation: you predict the result, change the controls, run the experiment and draw a conclusion. Inside are missions, theory, a quiz and a guided tour in which the camera moves and explains the key points.

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