Iron(III) oxide

This article is about a red-colored oxide of iron. For other uses, see Red Iron.

Iron(III) oxide or ferric oxide is the inorganic compound with the formula Fe₂O₃. It occurs in nature as the mineral hematite, which serves as the primary source of iron for the steel industry. It is also known as red iron oxide, especially when used in pigments.

Iron(III) oxide

Fe   O
Names
IUPAC name Iron(III) oxide
Other names ferric oxide, haematite, ferric iron, red iron oxide, rouge, maghemite, colcothar, iron sesquioxide, rust, ochre
Identifiers
CAS Number	1309-37-1 Y
3D model (JSmol)	Interactive image
ChEBI	CHEBI:50819 Y
ChemSpider	14147 N
ECHA InfoCard	100.013.790
EC Number	215-168-2
E number	E172(ii) (colours)
Gmelin Reference	11092
KEGG	C19424 N
PubChem CID	518696
RTECS number	NO7400000
UNII	1K09F3G675 Y
CompTox Dashboard (EPA)	DTXSID0029632
InChI InChI=1S/2Fe.3O Y[inchi] Key: JEIPFZHSYJVQDO-UHFFFAOYSA-N Y[inchi] InChI=1/2Fe.3O/rFe2O3/c3-1-4-2(3)5-1 Key: JEIPFZHSYJVQDO-ZVGCCQCPAC
SMILES O1[Fe]2O[Fe]1O2
Properties
Chemical formula	Fe2O3
Molar mass	159.687 g·mol−1
Appearance	Red solid
Odor	Odorless
Density	5.25 g/cm3[1]
Melting point	1,539 °C (2,802 °F; 1,812 K)[1] decomposes 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[2]
Solubility in water	Insoluble
Solubility	Soluble in diluted acids,[1] barely soluble in sugar solution[2] Trihydrate slightly soluble in aq. tartaric acid, citric acid, acetic acid[2]
Magnetic susceptibility (χ)	+3586.0x10−6 cm3/mol
Refractive index (nD)	n1 = 2.91, n2 = 3.19 (α, hematite)[3]
Structure
Crystal structure	Rhombohedral, hR30 (α-form)[4] Cubic bixbyite, cI80 (β-form) Cubic spinel (γ-form) Orthorhombic (ε-form)[5]
Space group	R3c, No. 161 (α-form)[4] Ia3, No. 206 (β-form) Pna21, No. 33 (ε-form)[5]
Point group	3m (α-form)[4] 2/m 3 (β-form) mm2 (ε-form)[5]
Coordination geometry	Octahedral (Fe3+, α-form, β-form)[4]
Thermochemistry[6]
Heat capacity (C)	103.9 J/mol·K[6]
Std molar entropy (S⦵298)	87.4 J/mol·K[6]
Std enthalpy of formation (ΔfH⦵298)	−824.2 kJ/mol[6]
Gibbs free energy (ΔfG⦵)	−742.2 kJ/mol[6]
Hazards
GHS labelling:
Pictograms	[7]
Signal word	Warning
Hazard statements	H315, H319, H335[7]
Precautionary statements	P261, P305+P351+P338[7]
NFPA 704 (fire diamond)	[8] 0 0 0
Threshold limit value (TLV)	5 mg/m3[1] (TWA)
Lethal dose or concentration (LD, LC):
LD50 (median dose)	10 g/kg (rats, oral)[8]
NIOSH (US health exposure limits):
PEL (Permissible)	TWA 10 mg/m3[9]
REL (Recommended)	TWA 5 mg/m3[9]
IDLH (Immediate danger)	2500 mg/m3[9]
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 verify (what is YN ?) Infobox references

Iron(III) oxide in a vial

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.

Iron(III) oxide is often called rust, since rust shares several properties and has a similar composition; however, in chemistry, rust is considered an ill-defined material, described as hydrous ferric oxide.^[10]

Ferric oxide is readily attacked by even weak acids. It is a weak oxidising agent, most famously when reduced by aluminium in the thermite reaction.

Structure

Fe₂O₃ can be obtained in various polymorphs. In the primary polymorph, α, iron adopts octahedral coordination geometry. That is, each Fe center is bound to six oxygen ligands. In the γ polymorph, some of the Fe sit on tetrahedral sites, with four 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.^[11] 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 converted from the alpha phase at high temperatures. It occurs naturally as the mineral maghemite. It is ferromagnetic and finds application in recording tapes,^[12] although ultrafine particles smaller than 10 nanometers are superparamagnetic. It can be prepared by thermal dehydratation of gamma iron(III) oxide-hydroxide. Another method involves the careful oxidation of iron(II,III) oxide (Fe₃O₄).^[12] The ultrafine particles can be prepared by thermal decomposition of iron(III) oxalate.

Other solid phases

Several other phases have been identified or claimed. The beta phase (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,^{[clarification needed]} pyrolysis of iron(III) chloride solution, or thermal decomposition of iron(III) sulfate.^[13]

The epsilon (ε) phase is rhombic, and shows properties intermediate between alpha and gamma, and may have useful magnetic properties applicable for purposes such as high density recording media for big data storage.^[14] Preparation of the pure epsilon phase has proven very challenging. Material with a high proportion of epsilon phase can be prepared by thermal transformation of the gamma phase. The epsilon phase is also metastable, transforming to the alpha phase at between 500 and 750 °C (930 and 1,380 °F). It can also be prepared by oxidation of iron in an electric arc or by sol-gel precipitation from iron(III) nitrate.^{[citation needed]} 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.^{[15][non-primary source needed]}

Additionally, at high pressure an amorphous form is claimed.^{[5][non-primary source needed]}

Liquid phase

Molten Fe₂O₃ is expected to have a coordination number of close to 5 oxygen atoms about each iron atom, based on measurements of slightly oxygen deficient supercooled liquid iron oxide droplets, where supercooling circumvents the need for the high oxygen pressures required above the melting point to maintain stoichiometry.^[16]

Hydrated iron(III) oxides

Several hydrates of Iron(III) oxide exist. 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 (α-Fe(O)OH) 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 (Fe^IIFe^III₂O₄), which is known as the mineral magnetite.

Fe(O)OH is soluble in acids, giving [Fe(H₂O)₆]³⁺. In concentrated aqueous alkali, Fe₂O₃ gives [Fe(OH)₆]³⁻.^[12]

Reactions

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.^[17]

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 produces magnetite, a black magnetic material that contains both Fe(III) and Fe(II):^[18]

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 chelating agents such as EDTA and oxalic acid.

Heating iron(III) oxides with other metal oxides or carbonates yields materials known as ferrates (ferrate (III)):^[18]

ZnO + Fe₂O₃ → Zn(FeO₂)₂

Preparation

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 FeO(OH), dehydrates around 200 °C.^[18][19]

2 FeO(OH) → Fe₂O₃ + H₂O

Uses

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.^[19]

Polishing

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.

Pigment

Sample of the red α- and yellow β-phases of hydrated of iron(III) oxide;^[2] both are useful as pigments.

Iron(III) oxide is also used as a pigment, under names "Pigment Brown 6", "Pigment Brown 7", and "Pigment Red 101".^[20] Some of them, e.g., Pigment Red 101 and Pigment Brown 6, are approved by the US Food and Drug Administration (FDA) for use in cosmetics. Iron oxides are used as pigments in dental composites alongside titanium oxides.^[21]

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). Its use in computer disks was superseded by cobalt alloy, enabling thinner magnetic films with higher storage density.^[22]

Photocatalysis

α-Fe₂O₃ has been studied as a photoanode for solar water oxidation.^[23] However, its efficacy is limited by a short diffusion length (2–4 nm) of photo-excited charge carriers^[24] and subsequent fast recombination, requiring a large overpotential to drive the reaction.^[25] Research has been focused on improving the water oxidation performance of Fe₂O₃ using nanostructuring,^[23] surface functionalization,^[26] or by employing alternate crystal phases such as β-Fe₂O₃.^[27]

Medicine

Calamine lotion, used to treat mild itchiness, is chiefly composed of a combination of zinc oxide, acting as astringent, and about 0.5% iron(III) oxide, the product's active ingredient, acting as antipruritic. The red color of iron(III) oxide is also mainly responsible for the lotion's pink color.

See also

• Chalcanthum

References

1. Haynes, p. 4.69
2. Comey, Arthur Messinger; Hahn, Dorothy A. (February 1921). A Dictionary of Chemical Solubilities: Inorganic (2nd ed.). New York: The MacMillan Company. p. 433.
3. Haynes, p. 4.141
4. 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. ISBN 978-1-118-07191-5.
5. 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). Univerzity Palackého. Retrieved 12 July 2014.
6. Haynes, p. 5.12
7. Sigma-Aldrich Co., Iron(III) oxide. Retrieved on 2014-07-12.
8. "SDS of Iron(III) oxide" (PDF). KJLC. England: Kurt J Lesker Company Ltd. 5 January 2012. Retrieved 12 July 2014.
9. NIOSH Pocket Guide to Chemical Hazards. "#0344". National Institute for Occupational Safety and Health (NIOSH).
10. PubChem. "Iron oxide (Fe2O3), hydrate". pubchem.ncbi.nlm.nih.gov. Retrieved 11 November 2020.
11. Greedan, J. E. (1994). "Magnetic oxides". In King, R. Bruce (ed.). Encyclopedia of Inorganic chemistry. New York: John Wiley & Sons. ISBN 978-0-471-93620-6.
12. Housecroft, Catherine E.; Sharpe, Alan G. (2008). "Chapter 22: d-block metal chemistry: the first row elements". Inorganic Chemistry (3rd ed.). Pearson. p. 716. ISBN 978-0-13-175553-6.
13. "Mechanism of Oxidation & Thermal Decomposition of Iron Sulphides" (PDF).
14. Tokoro, Hiroko; Namai, Asuka; Ohkoshi, Shin-Ichi (2021). "Advances in magnetic films of epsilon-iron oxide toward next-generation high-density recording media". Dalton Transactions. 50 (2). Royal Society of Chemistry: 452–459. doi:10.1039/D0DT03460F. PMID 33393552. S2CID 230482821. Retrieved 25 January 2021.
15. Dejoie, Catherine; Sciau, Philippe; Li, Weidong; Noé, Laure; Mehta, Apurva; Chen, Kai; Luo, Hongjie; Kunz, Martin; Tamura, Nobumichi; Liu, Zhi (2015). "Learning from the past: Rare ε-Fe₂O₃ in the ancient black-glazed Jian (Tenmoku) wares". Scientific Reports. 4: 4941. doi:10.1038/srep04941. PMC 4018809. PMID 24820819.
16. Shi, Caijuan; Alderman, Oliver; Tamalonis, Anthony; Weber, Richard; You, Jinglin; Benmore, Chris (2020). "Redox-structure dependence of molten iron oxides". Communications Materials. 1 (1): 80. Bibcode:2020CoMat...1...80S. doi:10.1038/s43246-020-00080-4.
17. Adlam; Price (1945). Higher School Certificate Inorganic Chemistry. Leslie Slater Price.
18. Handbook of Preparative Inorganic Chemistry, 2nd Ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 1661.
19. Greenwood, N. N.; Earnshaw, A. (1997). Chemistry of the Element (2nd ed.). Oxford: Butterworth-Heinemann. ISBN 978-0-7506-3365-9.
20. Paint and Surface Coatings: Theory and Practice. William Andrew Inc. 1999. ISBN 978-1-884207-73-0.
21. Banerjee, Avijit (2011). Pickard's Manual of Operative Dentistry. United States: Oxford University Press Inc., New York. p. 89. ISBN 978-0-19-957915-0.
22. Piramanayagam, S. N. (2007). "Perpendicular recording media for hard disk drives". Journal of Applied Physics. 102 (1): 011301–011301–22. Bibcode:2007JAP...102a1301P. doi:10.1063/1.2750414.
23. Kay, A.; Cesar, I.; Grätzel, M. (2006). "New Benchmark for Water Photooxidation by Nanostructured α-Fe₂O₃ Films". Journal of the American Chemical Society. 128 (49): 15714–15721. doi:10.1021/ja064380l. PMID 17147381.
24. Kennedy, J.H.; Frese, K.W. (1978). "Photooxidation of Water at α-Fe₂O₃ Electrodes". Journal of the Electrochemical Society. 125 (5): 709. Bibcode:1978JElS..125..709K. doi:10.1149/1.2131532.
25. Le Formal, F. (2014). "Back Electron–Hole Recombination in Hematite Photoanodes for Water Splitting". Journal of the American Chemical Society. 136 (6): 2564–2574. doi:10.1021/ja412058x. PMID 24437340.
26. Zhong, D.K.; Gamelin, D.R. (2010). "Photoelectrochemical Water Oxidation by Cobalt Catalyst ("Co−Pi")/α-Fe₂O₃ Composite Photoanodes: Oxygen Evolution and Resolution of a Kinetic Bottleneck". Journal of the American Chemical Society. 132 (12): 4202–4207. doi:10.1021/ja908730h. PMID 20201513.
27. Emery, J.D. (2014). "Atomic Layer Deposition of Metastable β-Fe₂O₃ via Isomorphic Epitaxy for Photoassisted Water Oxidation". ACS Applied Materials & Interfaces. 6 (24): 21894–21900. doi:10.1021/am507065y. OSTI 1355777. PMID 25490778.

Cited sources

• Haynes, William M., ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). CRC Press. ISBN 978-1439855119.

External links

• NIOSH Pocket Guide to Chemical Hazards