Iron(II,III) oxide

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Iron(II,III) oxide
Magnetite structure.jpg
IUPAC name
iron(II) iron(III) oxide
Other names
ferrous ferric oxide, ferroso ferric oxide, iron(II,III) oxide, magnetite, black iron oxide, lodestone, rust, iron(II) diiron(III) oxide
3D model (Jmol)
ECHA InfoCard 100.013.889


Molar mass 231.533 g/mol
Appearance solid black powder
Density 5 g/cm3
Melting point 1,597 °C (2,907 °F; 1,870 K)
2.42 [1]
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

Iron(II,III) oxide is the chemical compound with formula Fe3O4. It occurs in nature as the mineral magnetite. It is one of a number of iron oxides, the others being iron(II) oxide (FeO), which is rare, and iron(III) oxide (Fe2O3) also known as hematite. It contains both Fe2+ and Fe3+ ions and is sometimes formulated as FeO ∙ Fe2O3. This iron oxide is encountered in the laboratory as a black powder. It exhibits permanent magnetism and is ferrimagnetic, but is sometimes incorrectly described as ferromagnetic.[2] Its most extensive use is as a black pigment which is synthesised rather than being extracted from the naturally occurring mineral as the particle size and shape can be varied by the method of production.[3]


Under anaerobic conditions, ferrous hydroxide (Fe(OH)2) can be oxidized by water to form magnetite and molecular hydrogen. This process is described by the Schikorr reaction:

The well-crystallized magnetite (Fe3O4) is thermodynamically more stable than the ferrous hydroxide (Fe(OH)2 ).[4]

Magnetite can be prepared in the laboratory as a ferrofluid in the Massart method by mixing iron(II) chloride and iron(III) chloride in the presence of sodium hydroxide.[5] Magnetite can also be prepared by the chemical co-precipitation in presence of ammonia, which consist in a mixture of a solution 0.1 M of FeCl3·6H2O and FeCl2·4H2O with mechanic agitation of about 2000 rpm. The molar ratio of FeCl3:FeCl2 can be 2:1; heating this solution at 70 °C, and immediately the speed is elevated to 7500 rpm and adding quickly a solution of NH4OH (10 volume %), immediately a dark precipitate will be formed, which consists of nanoparticles of magnetite.[6] In both cases, the precipitation reaction rely on a quick transformation of acidic hydrolyzed iron ions into the spinel iron oxide structure, by hydrolysis at elevated pH values (above ca. 10).

Considerable efforts has been devoted towards controlling the particle formation process of magnetite nanoparticles due to the challenging and complex chemistry reactions involved in the phase transformations prior to the formation of the magnetite spinel structure.[7] Magnetite particles are of interests in bioscience applications such as in magnetic resonance imaging (MRI) since iron oxide magnetite nanoparticles represent a non-toxic alternative to currently employed gadolinium-based contrast agents. However, due to lack of control over the specific transformations involved in the formation of the particles, truly superparamagnetic particles have not yet been prepared from magnetite, i.e. magnetite nanoparticles that completely lose their permanent magnetic characteristic in the absence of an external magnetic field (which by definition show a coercivity of 0 A/m). The smallest values currently reported for nanosized magnetite particles is Hc = 8.5 A m−1,[8] whereas the largest reported magnetization value is 87 Am2 kg−1 for synthetic magnetite.[9][10]

Pigment quality Fe3O4, so called synthetic magnetite, can be prepared using processes that utilise industrial wastes, scrap iron or solutions containing iron salts (e.g. those produced as by-products in industrial processes such as the acid vat treatment (pickling) of steel):

  • Oxidation of Fe metal in the Laux process where nitrobenzene is treated with iron metal using FeCl2 as a catalyst to produce aniline:[3]
C6H5NO2 + 3 Fe + 2 H2O → C6H5NH2 + Fe3O4
  • Oxidation of FeII compounds, e.g. the precipitation of iron(II) salts as hydroxides followed by oxidation by aeration where careful control of the pH determines the oxide produced.[3]

Reduction of Fe2O3 with hydrogen:[11][12]

3Fe2O3 + H2 → 2Fe3O4 +H2O

Reduction of Fe2O3 with CO:[13]

3Fe2O3 + CO → 2Fe3O4 + CO2

Production of nano-particles can be performed chemically by taking for example mixtures of FeII and FeIII salts and mixing them with alkali to precipitate colloidal Fe3O4. The reaction conditions are critical to the process and determine the particle size.[14]


Reduction of magnetite ore by CO in a blast furnace is used to produce iron as part of steel production process:[2]

Controlled oxidation of Fe3O4 is used to produce brown pigment quality γ-Fe2O3 (maghemite):[15]

More vigorous calcining (roasting in air) gives red pigment quality α-Fe2O3 (hematite):[15]


Fe3O4 has a cubic inverse spinel structure which consists of a cubic close packed array of oxide ions where all of the Fe2+ ions occupy half of the octahedral sites and the Fe3+ are split evenly across the remaining octahedral sites and the tetrahedral sites.

Both FeO and γ-Fe2O3 have a similar cubic close packed array of oxide ions and this accounts for the ready interchangeability between the three compounds on oxidation and reduction as these reactions entail a relatively small change to the overall structure.[2] Fe3O4 samples can be non-stoichiometric.[2]

The ferrimagnetism of Fe3O4 arises because the electron spins of the FeII and FeIII ions in the octahedral sites are coupled and the spins of the FeIII ions in the tetrahedral sites are coupled but anti-parallel to the former. The net effect is that the magnetic contributions of both sets are not balanced and there is a permanent magnetism.[2]


Sample of magnetite, naturally occurring Fe3O4.

Fe3O4 is ferrimagnetic with a Curie temperature of 858 K. There is a phase transition at 120K, the so-called Verwey transition where there is a discontinuity in the structure, conductivity and magnetic properties.[16] This effect has been extensively investigated and whilst various explanations have been proposed, it does not appear to be fully understood.[17]

Fe3O4 is an electrical conductor with a conductivity significantly higher (X 106) than Fe2O3, and this is ascribed to electron exchange between the FeII and FeIII centres.[2]


Fe3O4 is used as a black pigment and is known as C.I pigment black 11 (C.I. No.77499).[15]

Fe3O4 is used as a catalyst in the Haber process and in the water gas shift reaction.[18] The latter uses an HTS (high temperature shift catalyst) of iron oxide stabilised by chromium oxide.[18] This iron-chrome catalyst is reduced at reactor start up to generate Fe3O4 from α-Fe2O3 and Cr2O3 to CrO3.[18]

Nano particles of Fe3O4 are used as contrast agents in MRI scanning.[19]

Ferumoxytol, also known as Feraheme and Rienso,[20] is an intravenous Fe3O4 preparation for treatment of anemia resulting from chronic kidney disease.[21] Ferumoxytol is manufactured and globally distributed by AMAG Pharmaceuticals.[20]

Along with sulfur and aluminium, it is an ingredient in a specific type of thermite useful for cutting steel.

Bluing is a passivation process that produces a layer of Fe3O4 on the surface of steel to protect it from rust.

Biological occurrence[edit]

Magnetite has been found as nano-crystals in magnetotactic bacteria (42-45 nm)[3] and in homing pigeon beak tissue[22]

See also[edit]


  1. ^ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
  2. ^ a b c d e f Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9. 
  3. ^ a b c d Rochelle M. Cornell, Udo Schwertmann 2007 The Iron Oxides: Structure, Properties, Reactions, Occurrences and Uses Wiley-VCH ISBN 3-527-60644-0
  4. ^ Ma, Ming; Zhang, Yu; Guo, Zhirui; Gu, Ning (2013). "Facile synthesis of ultrathin magnetic iron oxide nanoplates by Schikorr reaction". Nanoscale Research Letters. 8 (1): 16. doi:10.1186/1556-276X-8-16. 
  5. ^ Massart, R. (1981). "Preparation of aqueous magnetic liquids in alkaline and acidic media". IEEE Transactions on Magnetics. 17 (2): 1247–1248. doi:10.1109/TMAG.1981.1061188. 
  6. ^ Keshavarz, Sahar; Xu, Yaolin; Hrdy, Spencer; Lemley, Clay; Mewes, Tim; Bao, Yuping (2010). "Relaxation of Polymer Coated Fe3O4 Magnetic Nanoparticles in Aqueous Solution". IEEE Transactions on Magnetics. 46 (6): 1541–1543. doi:10.1109/TMAG.2010.2040588. 
  7. ^ Jolivet, Jean-Pierre; Chanéac, Corinne; Tronc, Elisabeth (2004). "Iron oxide chemistry. From molecular clusters to extended solid networks". Chemical Communications (5): 477–483. doi:10.1039/B304532N. 
  8. ^ Ström, Valter; Olsson, Richard T.; Rao, K. V. (2010). "Real-time monitoring of the evolution of magnetism during precipitation of superparamagnetic nanoparticles for bioscience applications". Journal of Materials Chemistry. 20 (20): 4168. doi:10.1039/C0JM00043D. 
  9. ^ Fang, Mei; Ström, Valter; Olsson, Richard T.; Belova, Lyubov; Rao, K. V. (2011). "Rapid mixing: A route to synthesize magnetite nanoparticles with high moment". Applied Physics Letters. 99 (22): 222501. doi:10.1063/1.3662965. 
  10. ^ Fang, Mei; Ström, Valter; Olsson, Richard T; Belova, Lyubov; Rao, K. V. (2012). "Particle size and magnetic properties dependence on growth temperature for rapid mixed co-precipitated magnetite nanoparticles". Nanotechnology. 23 (14): 145601. doi:10.1088/0957-4484/23/14/145601. 
  11. ^ US patent 2596954, 1947, Process for reduction of iron ore to magnetiteHeath T.D.
  12. ^ A. Pineau; N. Kanari; I. Gaballah (2006). "Kinetics of reduction of iron oxides by H2 Part I: Low temperature reduction of hematite". Thermochimica Acta. 447 (1): 89–100. doi:10.1016/j.tca.2005.10.004. 
  13. ^ Hayes P. C.; Grieveson P. (1981). "The effects of nucleation and growth on the reduction of Fe2O3 to Fe3O4". Metallurgical and Materials Transactions B. 12 (2): 319–326. doi:10.1007/BF02654465. 
  14. ^ Arthur T. Hubbard (2002) Encyclopedia of Surface and Colloid Science CRC Press, ISBN 0-8247-0796-6
  15. ^ a b c Gunter Buxbaum, Gerhard Pfaff (2005) Industrial Inorganic Pigments 3d edition Wiley-VCH ISBN 3-527-30363-4
  16. ^ Verwey E. J. W. (1939). "Electronic Conduction of Magnetite (Fe3O4) and its Transition Point at Low Temperatures". Nature. 144 (3642): 327–328 (1939). doi:10.1038/144327b0. 
  17. ^ Walz, F. (2002). "The Verwey transition - a topical review". Journal of Physics: Condensed Matter. 14 (12): R285–R340. doi:10.1088/0953-8984/14/12/203. 
  18. ^ a b c Sunggyu Lee (2006) Encyclopedia of Chemical Processing CRC Press ISBN 0-8247-5563-4
  19. ^ Babes L; Denizot B; Tanguy G; Le Jeune J.J.; Jallet P. (1999). "Synthesis of Iron Oxide Nanoparticles Used as MRI Contrast Agents: A Parametric Study". Journal of Colloid and Interface Science. 212 (2): 474–482. doi:10.1006/jcis.1998.6053. PMID 10092379. 
  20. ^ a b "AMAG Pharmaceuticals and Takeda Announce Mutual Termination of Agreement to License, Develop and Commercialize Ferumoxytol in Ex-U.S. Territories, Including Europe" (Press release). AMAG Pharmaceuticals. 29 December 2014 – via Seeking Alpha. 
  21. ^ Schwenk, Michael H. (2010). "Ferumoxytol: A New Intravenous Iron Preparation for the Treatment of Iron Deficiency Anemia in Patients with Chronic Kidney Disease". Pharmacother. 30 (1): 70–9 – via Medscape. (registration required)
  22. ^ Hanzlik M.; Heunemann C.; Holtkamp-Rötzler E.; Winklhofer M.; Petersen N.; Fleissner G (2000). "Superparamagnetic Magnetite in the Upper Beak Tissue of Homing Pigeons". BioMetals. 13 (4): 325–331. doi:10.1023/A:1009214526685. PMID 11247039.