Iron (III) compounds

Iron. In this state, iron has five d electrons, but does not show any strong resemblance to manganese(II), except that most iron(III) compounds show high paramagnetism, i.e. its outermost shell electrons remain unpaired.

Iron(III) chloride is a black, essentially covalent solid, in which each iron atom is surrounded octahedrally by six chlorine atoms. It is prepared by direct combination of iron with chlorine or by dehydration of the hydrated chloride.

When the anhydrous solid is heated, it vaporises to form first Fe₂Cl₆ molecules, then the monomer FeCl₃ and finally FeCl₂ and chlorine. It fumes in air (with hydrolysis) and dissolves readily in water to give a yellow (dilute) or brown (concentrated) solution, which is strongly acidic. Crystallisation gives the yellow hydrate FeCl₃^.6H₂O which has the structure [FeCl₂(H₂O)₄]Cl₂H₂O, i.e. contains the octahedral complex ion [FeCl₂(H₂O)₄]⁺; ions of this general type are responsible for the colours of the aqueous solution of iron(III) chloride. In the presence of excess chloride ion, both tetrahedral [FeCl₄]^- and octahedral [FeCl₆]^3- can be formed.

Iron(III) chloride forms numerous addition compounds, especially with organic molecules which contain donor atoms, for example ethers, alcohols, aldehydes, ketones and amines. Anhydrous iron(III) chloride is soluble in, for example, ether, and can be extracted into this solvent from water; the extraction is more effective in the presence of chloride ion.

Of other iron(III) halides, iron(III) bromide and iron(III) iodide decompose rather readily into the +2 halide and halogen.

Iron(III) sulfate Fe₂(SO₄)₃ forms very hygroscopic white crystals. It forms the crystal hydrate Fe₂(SO₄)₃·9H₂O (yellow crystals). Iron(III) sulfate is greatly hydrolyzed In aqueous solutions, it forms double salts—alums— with alkali metal and ammonium sulfates. An example isammonium iron alum (NH₄)Fe(SO₄)₂·12H₂O—light violet crystals well soluble in water. When roasted to above 500 °C, iron(III) sulfate decomposes as follows:

Fe₂(SO₄)₃ = Fe₂O₃ + 3SO₃↑.

Iron(III) sulfate is employed, like FeCl₃, as a coagulant in water purification, and also for pickling metals. A solution of Fe₂(SO₄)₃ can dissolve Cu₂S and CuS to form copper(II) sulfate; this property is used in the hydrometallurgical preparation of copper.

When alkalis react with solutions of iron(III) salts, red-brown iron (III) hydroxide Fe(OH)₃ precipitates. It is insoluble in an excess of the alkali.

Iron(III) oxides and hydroxide. If an aqueous solution of an iron(III) salt is treated with alkali, a red-brown precipitate of Iron(III) hydroxide' is obtained; this is probably best represented as FeO(OH). At strong heating it gives the red oxide Fe₂O₃. Iron(III) oxide occurs naturally as haematite, and can also be prepared by strong heating of iron(II) sulfate:

2FeSO₄ = Fe₂O₃ + SO₂ +SO₃,

It shows some amphoteric behaviour, since it dissolves in alkali (concentrated aqueous or fused) to give a ferrate(III); the equation may be written as

Fe₂O₃ + 2OH^- = 2FeO₂^- + H₂O

Iron(II) oxide exists in two forms, the red a-form (paramagnetic) and the y-form (ferromagnetic) obtained by careful heating of FeO(OH). The a-form is used as a red pigment, as a metal polish ("jeweller's rouge') and as a catalyst.

The mixed oxide Fe₃O₄ (triiron tetroxide) is a black solid, which occurs naturally as magnetite; it is formed when iron(III) oxide is strongly heated, and its structure is effectively made up of oxide (O^2-) and iron(II) and iron(III) ions.

Ferrites. When iron(III) oxide is fused with sodium or potassium carbonates, ferrites are formed—salts of ferrous acid HFeO₂ that has not been obtained in the free state, for instance sodium ferrite NaFeO₂:

Fe₂O₃ + Na₂CO₃ = 2NaFeO₂+ C0₂↑.

In engineering, the name ferrites or ferrite materials is applied to the products obtained in sintering powders of iron(III) oxide and oxides of some divalent metals, for instance, nickel, zinc, and manganese. Sintering is conducted at 1000 to 1400 °C. Ferrites have valuable magnetic properties and high electrical resistance, which underlies the small value of the electrical losses in them. Ferrites find broad application in communication equipment, com­puters, and in automatic and remote, control equipment.

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