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Ferric oxide

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Ferric oxide

Iron(III) oxide
CAS number 1309-37-1 YesY
PubChem 518696
ChemSpider 14147 N
UNII 1K09F3G675 YesY
KEGG C19424 YesY
ChEBI CHEBI:50819 YesY
RTECS number NO7400000
Jmol-3D images Image 1
Molecular formula Fe2O3
Molar mass 159.69 g/mol
Appearance red-brown solid
Odor odorless
Density 5.242 g/cm3, solid
Melting point

1566 °C, 1839 K, 2851 °F (decomposes)

Solubility in water insoluble
Crystal structure rhombohedral
Std enthalpy of
−826 kJ·mol−1[1]
Standard molar
90 J·mol−1·K−1[1]
EU classification not listed
Flash point non-flammable
Related compounds
Other anions iron(III) fluoride
Other cations manganese(III) oxide, cobalt(III) oxide
Related compounds iron(II) oxide, iron(II,III) oxide
 N (verify) (what is: YesY/N?)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Iron(III) oxide or ferric oxide is the inorganic compound with the formula Fe2O3. It is one of the three main oxides of iron, the other two being iron(II) oxide (FeO), which is rare, and iron(II,III) oxide (Fe3O4), which also occurs naturally as the mineral magnetite. As the mineral known as hematite, Fe2O3 is the main source of the iron for the steel industry. Fe2O3 is ferromagnetic, dark red, and readily attacked by acids. Iron(III) oxide is often called rust, and to some extent this label is useful, because rust shares several properties and has a similar composition. To a chemist, rust is considered an ill-defined material, described as hydrated ferric oxide.


Fe2O3 can be obtained in various polymorphs. In the main ones, α and γ, iron adopts octahedral coordination geometry. That is, each Fe center is bound to six oxygen ligands.

Alpha phase

α-Fe2O3 has the rhombohedral, corundum (α-Al2O3) structure and is the most common form. It occurs naturally as the mineral hematite which is mined as the main ore of iron. It is antiferromagnetic below ~260 K (Morin transition temperature), and exhibits weak ferromagnetism between 260 K and the Néel temperature, 950 K.[2] It is easy to prepare using both thermal decomposition and precipitation in the liquid phase. Its magnetic properties are dependent on many factors, e.g. pressure, particle size, and magnetic field intensity.

Gamma phase

γ-Fe2O3 has a cubic structure. It is metastable and converts to the alpha phase at high temperatures. It occurs naturally as the mineral maghemite. It is ferromagnetic and finds application in recording tapes,[3] although ultrafine particles smaller than 10 nanometers are superparamagnetic. It can be prepared by thermal dehydratation of gamma iron(III) oxide-hydroxide, careful oxidation of iron(II,III) oxide. Another method involves the careful oxidation of Fe3O4.[3] The ultrafine particles can be prepared by thermal decomposition of iron(III) oxalate.

Other phases

Several other phases have been identified or claimed. The beta-phase is cubic body centered (space group Ia3), metastable, and at temperatures above 500 °C (930 °F) converts to alpha phase. It can be prepared by reduction of hematite by carbon, pyrolysis of iron(III) chloride solution, or thermal decomposition of iron(III) sulfate. The epsilon phase is rhombic, and shows properties intermediate between alpha and gamma. So far has not been prepared in pure form; it is always mixed with the alpha phase or gamma phases. Material with a high proportion of epsilon phase can be prepared by thermal transformation of the gamma phase. This phase is also metastable, transforming to the alpha phase at between 500 and 750 °C (930 and 1,380 °F). Can also be prepared by oxidation of iron in an electric arc or by sol-gel precipitation from iron(III) nitrate. Additionally at high pressure an amorphous form is claimed.[4]

Hydrated iron(III) oxides

Several hydrates of Iron(III) oxide exists. When alkali is added to solutions of soluble Fe(III) salts, a red-brown gelatinous precipitate forms. This is not Fe(OH)3, but Fe2O3·H2O (also written as Fe(O)OH). Several forms of the hydrated oxide of Fe(III) exist as well. The red lepidocrocite γ-Fe(O)OH, occurs on the outside of rusticles, and the orange goethite, which occurs internally in rusticles. When Fe2O3·H2O is heated, it loses its water of hydration. Further heating at 1670 K converts Fe2O3 to black Fe3O4 (FeIIFeIII2O4), which is known as the mineral magnetite. Fe(O)OH is soluble in acids, giving [Fe(OH2)6]3+. In concentrated aqueous alkali, Fe2O3 gives [Fe(OH)6]3−.[3]


The most important reaction is its carbothermal reduction, which gives iron used in steel-making:

2 Fe2O3 + 3 C → 4 Fe + 3 CO2

Another redox reaction is the extremely exothermic thermite reaction with aluminium.[5]

2 Al + Fe2O3 → 2 Fe + Al2O3

This process is used to weld thick metals such as rails of train tracks by using a ceramic container to funnel the molten iron in between two sections of rail. Thermite is also used in weapons and making small-scale cast-iron sculptures and tools.

Partial reduction with hydrogen at about 400 °C gives magnetite, a black magnetic material that contains both Fe(III) and Fe(II):[6]

3 Fe2O3 + H2 → 2 Fe3O4 + H2O

Iron(III) oxide is insoluble in water [needs citatation] but dissolves readily in strong acid, e.g. hydrochloric and sulfuric acids. It also dissolves well in solutions of the chelating agents such as EDTA and oxalic acid.

Heating iron(III) oxides with other metal oxides or carbonates yields materials known as ferrates:[6]

ZnO + Fe2O3 → Zn(FeO2)2


Iron (III) oxide is a product of the oxidation of iron. It can be prepared in the laboratory by electrolyzing a solution of sodium bicarbonate, an inert electrolyte, with an iron anode:

4 Fe + 3 O2 + 2 H2O → 4 FeO(OH)

The resulting hydrated iron(III) oxide, written here as Fe(O)OH, dehydrates around 200 °C.[6][7]

2 FeO(OH) → Fe2O3 + H2O

It can also be prepared by the thermal decomposition of iron (III) hydroxide under temperature above 200 °C.

2 Fe(OH)3 → Fe2O3 + 3H2O


Iron industry

The overwhelming application of iron(III) oxide is as the feedstock of the steel and iron industries, e.g. the production of iron, steel, and many alloys.[7]


A very fine powder of ferric oxide is known as "jeweler's rouge", "red rouge", or simply rouge. It is used to put the final polish on metallic jewelry and lenses, and historically as a cosmetic. Rouge cuts more slowly than some modern polishes, such as cerium(IV) oxide, but is still used in optics fabrication and by jewelers for the superior finish it can produce. When polishing gold, the rouge slightly stains the gold, which contributes to the appearance of the finished piece. Rouge is sold as a powder, paste, laced on polishing cloths, or solid bar (with a wax or grease binder). Other polishing compounds are also often called "rouge", even when they do not contain iron oxide. Jewelers remove the residual rouge on jewelry by use of ultrasonic cleaning. Products sold as "stropping compound" are often applied to a leather strop to assist in getting a razor edge on knives, straight razors, or any other edged tool.


Iron(III) oxide is also used as a pigment, under names "Pigment Brown 6", "Pigment Brown 7", and "Pigment Red 101".[8] Some of them, e.g. Pigment Red 101 and Pigment Brown 6, are Food and Drug Administration (FDA)-approved for use in cosmetics. Iron oxides are used as pigments in dental composites alongside titanium oxides.[9]

Hematite is the characteristic component of the Swedish paint color Falu red.

Magnetic Recording

Iron(III) oxide is the most common magnetic particle used for magnetic recording in hard drives and audio tape. The oxide is mixed with a binder and coated onto a plastic film substrate. Magnetic tape is used in audio and video cassettes and reel to reel tape recording of all kinds. The oxide is magnetized in the recording process in the pattern of the audio signals. When the tape is played back, the now magnetized ferric oxide is run over the tape heads, generating an electrical signal that is fed to an audio amplifier and speakers to reproduce the recorded sound.


α-Fe2O3 has been studied as a photoanode for the water-splitting reaction for over 25 years.[10]

See also


External links

  • NIOSH Pocket Guide to Chemical Hazards

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