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Iron(III) oxide

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Title: Iron(III) oxide  
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Iron(III) oxide

Iron(III) oxide
Haematite unit cell
Sample of iron(III) oxide
IUPAC name
Iron(III) oxide
Other names
Ferric oxide, Hematite, Ferric iron, Red iron oxide, Rouge, Maghemite, Colcothar, Iron sesquioxide, Rust, Ochre
ChemSpider  N
EC number 215-168-2
Jmol-3D images Image
RTECS number NO7400000
Molar mass 159.69 g·mol−1
Appearance Red-brown solid
Odor Odorless
Density 5.242 g/cm3[1]
Melting point 1,539–1,565 °C (2,802–2,849 °F; 1,812–1,838 K)[2]
105 °C (221 °F; 378 K)
β-dihydrate, decomposes
150 °C (302 °F; 423 K)
β-monohydrate, decomposes
50 °C (122 °F; 323 K)
α-dihydrate, decomposes
92 °C (198 °F; 365 K)
α-monohydrate, decomposes[3]
Solubility Soluble in diluted acids,[2] sugar solution
Trihydrate slighty soluble in aq. tartaric acid, citric acid, CH3COOH[3]
Rhombohedral, hR30 (α-form)[4]
Cubic bixbyite, cI80 (β-form)
Cubic spinel (γ-form)
Orthorhombic (ε-form)[5]
R3c, No. 161 (α-form)[4]
Ia3, No. 206 (β-form)
Pna21, No. 33 (ε-form)[5]
3m (α-form)[4]
2/m 3 (β-form)
mm2 (ε-form)[5]
Octahedral (Fe3+, α-form, β-form)[4]
103.9 J/mol·K[2]
87.4 J/mol·K[2]
−824.2 kJ/mol[2]
−742.2 kJ/mol[2]
GHS pictograms The exclamation-mark pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS)[6]
GHS signal word Warning
H315, H319, H335[6]
P261, P305+351+338[6]
Irritant Xi
R-phrases R36/37/38
S-phrases S26
NFPA 704
5 mg/m3[2] (TWA)
Lethal dose or concentration (LD, LC):
LD50 (Median dose)
10 g/kg (rats, oral)[1]
US health exposure limits (NIOSH):
PEL (Permissible)
TWA 10 mg/m3[7]
REL (Recommended)
TWA 5 mg/m3[7]
2500 mg/m3[7]
Related compounds
Other anions
Iron(III) fluoride
Other cations
Manganese(III) oxide
Cobalt(III) oxide
Related iron oxides
Iron(II) oxide
Iron(II,III) oxide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
 N  (: Y/N?)

Iron(III) oxide or ferric oxide is the 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 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.


  • Structure 1
    • Alpha phase 1.1
    • Gamma phase 1.2
    • Other phases 1.3
  • Hydrated iron(III) oxides 2
  • Reactions 3
  • Preparation 4
  • Uses 5
    • Iron industry 5.1
    • Polishing 5.2
    • Pigment 5.3
    • Magnetic recording 5.4
    • Photocatalyst 5.5
    • Medicine 5.6
  • See also 6
  • References 7
  • External links 8


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.[8] 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 [9] The ultrafine particles can be prepared by thermal decomposition of iron(III) oxalate.

Other phases

Several other phases have been identified or claimed. The β-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, and may have useful magnetic properties. Preparation of the pure epsilon phase has proven very challenging due to contamination with alpha and 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.[5] Recent research has revealed epsilon iron(III) oxide in ancient Chinese Jian ceramic glazes, which may provide insight into ways to produce that form in the lab. [10]

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 [9]


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

Fe2O3 + 3 CO → 2 Fe + 3 CO2

Another redox reaction is the extremely exothermic thermite reaction with aluminium.[11]

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):[12]

3 Fe2O3 + H2 → 2 Fe3O4 + H2O

Iron(III) oxide is insoluble in water 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:[12]

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.[12][13]

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.[13]


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.


Two different colors at different hydrate phase (α = red, β = yellow) of iron(III) oxide hydrate[3] and they are useful as a pigment.

Iron(III) oxide is also used as a pigment, under names "Pigment Brown 6", "Pigment Brown 7", and "Pigment Red 101".[14] 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.[15]

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

Magnetic recording

Iron(III) oxide was the most common magnetic particle used in all types of magnetic storage and recording media, including magnetic disks (for data storage) and magnetic tape (used in audio and video recording as well as data storage). However, modern magnetic storage media - in particular, the hard disk drives - use more advanced thin film technology, which may consist of a stack of 15 layers or more. [16]


α-Fe2O3 has been studied as a photoanode for the water-splitting reaction for over 25 years.[17]


A mixture of zinc oxide with about 0.5% iron(III) oxide is called calamine, which is the active ingredient of calamine lotion.

See also


  1. ^ a b c d "SDS of Iron(III) oxide" (PDF). England: Kurt J Lesker Company Ltd. 2012-01-05. Retrieved 2014-07-12. 
  2. ^ a b c d e f g Lide, David R., ed. (2009).  
  3. ^ a b c Comey, Arthur Messinger; Hahn, Dorothy A. (1921-02). A Dictionary of Chemical Solubilities: Inorganic (2nd ed.). New York: The MacMillan Company. p. 433. 
  4. ^ a b c d Ling, Yichuan; Wheeler, Damon A.; Zhang, Jin Zhong; Li, Yat (2013). Zhai, Tianyou; Yao, Jiannian, eds. One-Dimensional Nanostructures: Principles and Applications. (Hoboken, New Jersey: John Wiley & Sons, Inc.). p. 167.  
  5. ^ a b c d Vujtek, Milan; Zboril, Radek; Kubinek, Roman; Mashlan, Miroslav. "Ultrafine Particles of Iron(III) Oxides by View of AFM – Novel Route for Study of Polymorphism in Nano-world" (PDF). Czech. Retrieved 2014-07-12. 
  6. ^ a b c Sigma-Aldrich Co., Iron(III) oxide. Retrieved on 2014-07-12.
  7. ^ a b c "NIOSH Pocket Guide to Chemical Hazards #0344".  
  8. ^ Greedon, J. E. (1994). "Magnetic oxides". In King, R. Bruce. Encyclopedia of Inorganic chemistry. New York: John Wiley & Sons.  
  9. ^ a b c .Housecroft, Catherine E.; Sharpe, Alan G. (2008). "Chapter 22: d-block metal chemistry: the first row elements". Inorganic Chemistry, 3rd Edition. Pearson. p. 716.  
  10. ^
  11. ^ Adlam; Price (1945). Higher School Certificate Inorganic Chemistry. Leslie Slater Price. 
  12. ^ a b c Handbook of Preparative Inorganic Chemistry, 2nd Ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 1661.
  13. ^ a b Greenwood, N. N.; Earnshaw, A. (1997). Chemistry of the Element (2nd ed.). Oxford: Butterworth-Heinemann.  
  14. ^ Paint and Surface Coatings: Theory and Practice. William Andrew Inc.  
  15. ^ Banerjee, Avijit (2011). Pickard's Manual of Operative Dentistry. United States: Oxford University Press Inc., New York. p. 89.  
  16. ^ S.N. Piramanayagam, J. Appl. Phys. 102, 011301 (2007).
  17. ^ Kay, A., Cesar, I. and Gratzel, M, Journal of the American Chemical Society 2006, 128, 15714-15721

External links

  • NIOSH Pocket Guide to Chemical Hazards
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