Hubbry Logo
Iron(II) oxideIron(II) oxideMain
Open search
Iron(II) oxide
Community hub
Iron(II) oxide
logo
8 pages, 0 posts
0 subscribers
Be the first to start a discussion here.
Be the first to start a discussion here.
Iron(II) oxide
Iron(II) oxide
Iron(II) oxide
View on Wikipedia
from Wikipedia
Iron(II) oxide
Iron(II) oxide
Iron(II) oxide
Names
IUPAC name
Iron(II) oxide
Other names
Ferrous oxide, Iron monoxide, Wüstite
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.014.292 Edit this at Wikidata
13590
UNII
  • InChI=1S/Fe.O checkY
    Key: UQSXHKLRYXJYBZ-UHFFFAOYSA-N checkY
  • InChI=1/Fe.O/rFeO/c1-2
    Key: UQSXHKLRYXJYBZ-WPTVXXAFAB
  • [Fe]=O
Properties
FeO
Molar mass 71.844 g/mol
Appearance black crystals
Density 5.745 g/cm3
Melting point 1,377 °C (2,511 °F; 1,650 K)[1]
Boiling point 3,414 °C (6,177 °F; 3,687 K)
Insoluble
Solubility insoluble in alkali, alcohol
dissolves in acid
+7200×10−6 cm3/mol
2.23
Hazards
Occupational safety and health (OHS/OSH):
Main hazards
 can be combustible under specific conditions[2]
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 1: Exposure would cause irritation but only minor residual injury. E.g. turpentineFlammability 1: Must be pre-heated before ignition can occur. Flash point over 93 °C (200 °F). E.g. canola oilInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
1
1
0
200 °C (392 °F; 473 K)
Safety data sheet (SDS) ICSC 0793
Related compounds
Other anions
 Iron(II) sulfide
Iron(II) selenide
Iron(II) telluride
Other cations
 Manganese(II) oxide
Cobalt(II) oxide
Related Iron oxides
 Iron(II,III) oxide
Iron(III) oxide
Related compounds
 Iron(II) fluoride
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
checkY verify (what is checkY☒N ?)

Iron(II) oxide or ferrous oxide is the inorganic compound with the formula FeO. Its mineral form is known as wüstite.[3][4] One of several iron oxides, it is a black-colored powder that is sometimes confused with rust, the latter of which consists of hydrated iron(III) oxide (ferric oxide). Iron(II) oxide also refers to a family of related non-stoichiometric compounds, which are typically iron deficient with compositions ranging from Fe0.84O to Fe0.95O.[5]

Preparation

[edit]

FeO can be prepared by the thermal decomposition of iron(II) oxalate.

FeC2O4 → FeO + CO2 + CO

The procedure is conducted under an inert atmosphere to avoid the formation of iron(III) oxide (Fe2O3). A similar procedure can also be used for the synthesis of manganous oxide and stannous oxide.[6][7]

Stoichiometric FeO can be prepared by heating Fe0.95O with metallic iron at 770 °C and 36 kbar.[8]

Reactions

[edit]

FeO is thermodynamically unstable below 575 °C, tending to disproportionate to metal and Fe3O4:[5]

4 FeO → Fe + Fe3O4

Structure

[edit]

Iron(II) oxide adopts the cubic, rock salt structure, where iron atoms are octahedrally coordinated by oxygen atoms and the oxygen atoms octahedrally coordinated by iron atoms. The non-stoichiometry occurs because of the ease of oxidation of FeII to FeIII effectively replacing a small portion of FeII with two-thirds their number of FeIII, which take up tetrahedral positions in the close packed oxide lattice.[8]

In contrast to the crystalline solid, in the molten state iron atoms are coordinated by predominantly 4 or 5 oxygen atoms.[9]

Below 200 K there is a minor change to the structure which changes the symmetry to rhombohedral and samples become antiferromagnetic.[8][10]

Occurrence in nature

[edit]

Iron(II) oxide makes up approximately 9% of the Earth's mantle. Within the mantle, it may be electrically conductive, which is a possible explanation for perturbations in Earth's rotation not accounted for by accepted models of the mantle's properties.[11]

Uses

[edit]

Iron(II) oxide is used as a pigment. It is FDA-approved for use in cosmetics and it is used in some tattoo inks. It can also be used as a phosphate remover from home aquaria.

See also

[edit]

References

[edit]
[edit]
Revisions and contributorsEdit on WikipediaRead on Wikipedia

Iron(II) oxide

View on Grokipedia
from Grokipedia
Iron(II) oxide, also known as ferrous oxide, is an with the FeO, consisting of iron in the +2 bonded to oxygen. It occurs naturally as the mineral wüstite (typically non-stoichiometric Fe_{1-x}O), a rare found in meteorites and certain terrestrial rocks, and is typically synthesized for industrial use as a black crystalline solid or powder. With a molecular weight of 71.84 g/mol, it has a of approximately 5.7–6.0 g/cm³, melts at 1377 °C, and is insoluble in but soluble in acids. Chemically, iron(II) oxide is unstable in air and readily oxidizes to iron(III) oxide, particularly when finely divided, making it pyrophoric and requiring inert handling conditions. It exhibits basic properties, absorbing carbon dioxide to form carbonates, and serves as a reducing agent in various reactions due to the Fe²⁺ cation. Preparation often involves thermal decomposition of iron(II) oxalate or reduction of higher iron oxides, such as heating Fe₂O₃ with hydrogen or carbon monoxide under controlled conditions. Iron(II) oxide finds applications in and manufacturing, particularly as a component in ceramics, for heat-absorbing properties, and enamel coatings. It is also used as a catalyst in chemical processes, a colorant in , and in , where it acts as an intermediate in formation to remove impurities.

Properties

Physical properties

Iron(II) oxide, also known as in its mineral form, appears as a black solid, typically in the form of a hygroscopic , pellets, or cubic . Its is 71.844 g/mol. The compound has a of 5.7 g/cm³. It melts at 1,377 °C (1,650 ) and boils at 3,414 °C (3,687 ). Iron(II) oxide is insoluble in and alkalies but soluble in dilute acids.
PropertyValueConditions/Source
Specific heat capacity49.93 J/mol·K298 K
Thermal conductivity1.8–2.5 W/m·K to 1,000 °C
Electrical resistivity~0.01 Ω·cmElevated temperatures (semiconducting behavior)
Optically, iron(II) oxide displays a black color across its forms and has a refractive index of 2.23.

Chemical properties

Iron(II) oxide, denoted stoichiometrically as FeO, consists of iron in the +2 oxidation state (Fe²⁺) bonded to the oxide ion (O²⁻) in an ionic lattice. As a basic metal oxide, FeO reacts with acids to form the corresponding iron(II) salts and water. For example, it dissolves in hydrochloric acid according to the reaction: FeO+2HCl→FeCl2+H2O\text{FeO} + 2\text{HCl} \rightarrow \text{FeCl}_2 + \text{H}_2\text{O} FeO exhibits hygroscopic behavior, absorbing moisture from the atmosphere, which facilitates its slow oxidation in air to higher iron oxides such as Fe₂O₃, resulting in a color change from black to reddish-brown. In terms of electronic properties, stoichiometric FeO displays antiferromagnetic ordering below its Néel temperature of approximately 198 K, arising from the antiferromagnetic coupling of Fe²⁺ ions in its rock-salt structure, which influences its reactivity in magnetic and electrochemical contexts. The solubility of FeO is pH-dependent; it is generally insoluble in neutral water but can hydrolyze to form iron(II) hydroxide (Fe(OH)₂), whose solubility increases at lower pH due to protonation and dissolution as aqueous Fe²⁺ species, while at higher pH values above 10, precipitation of Fe(OH)₂ dominates, limiting overall iron(II) solubility.

Structure

Crystal structure

Iron(II) oxide, in its ideal stoichiometric form, crystallizes in the cubic rock salt (NaCl-type) structure with the space group Fm3ˉmFm\bar{3}m (No. 225). In this arrangement, the Fe²⁺ and O²⁻ ions occupy the cation and anion sites of a face-centered cubic lattice, respectively, forming a three-dimensional network where alternating layers of iron and oxygen ions stack along the <111> directions. The unit cell contains four formula units of FeO, with the ions positioned at the corners and face centers of the cube. The lattice parameter aa for the ideal structure is 4.32 at ambient conditions, corresponding to Fe–O bond lengths of approximately 2.16 . This parameter reflects the balance between the ionic radii of Fe²⁺ (high-spin, 0.78 ) and (1.40 ), which dictate the close-packed geometry. Both cations and anions exhibit octahedral coordination, with each Fe²⁺ surrounded by six ions and vice versa, resulting in a of 6 for all ions and strong directional bonding along the edges of the octahedra. The bonding in iron(II) oxide is primarily ionic, driven by the electrostatic attraction between the divalent ions, but it possesses partial covalent character arising from the overlap of Fe 3d orbitals with O 2p orbitals. This hybridization contributes to the stability of the rock salt phase and influences electronic properties, such as the narrowing of the band gap compared to purely ionic analogs. In the molten state, above the melting point of approximately 1370 °C, the local structure transitions to a more disordered liquid with iron atoms coordinated by 4 to 5 oxygen atoms on average, as determined by empirical potential structure refinement of neutron diffraction data. While the ideal structure provides a foundational model, real samples of iron(II) oxide often exhibit non-stoichiometric deviations that alter the lattice slightly.

Defects and non-stoichiometry

Iron(II) oxide, commonly known as in its mineral form, exhibits significant non-stoichiometry due to inherent structural imperfections that deviate from the ideal 1:1 Fe:O ratio. This non-stoichiometric nature arises primarily from cation deficiencies in the rock salt lattice, leading to a composition best represented by the formula \ceFe1−xO\ce{Fe_{1-x}O}, where xx ranges from 0.05 to 0.16, corresponding to \ceFe0.84O\ce{Fe_{0.84}O} to \ceFe0.95O\ce{Fe_{0.95}O}. These defects are ubiquitous in synthesized or occurring under equilibrium conditions at high temperatures. The defect mechanism involves the formation of iron cation vacancies to accommodate the oxygen-rich composition, with charge neutrality maintained by the oxidation of surrounding iron ions. Specifically, for each cation vacancy created by the removal of an \ceFe2+\ce{Fe^{2+}} ion, two adjacent \ceFe2+\ce{Fe^{2+}} ions are oxidized to \ceFe3+\ce{Fe^{3+}} to balance the positive charge deficit, conceptually illustrated as 3\ceFe2+→2\ceFe3++V\ceFe3 \ce{Fe^{2+}} \rightarrow 2 \ce{Fe^{3+}} + V_{\ce{Fe}}. This results in a structure containing both divalent and trivalent iron cations clustered around vacancies, often forming defect complexes that influence long-range ordering. These structural defects profoundly affect the physical properties of . The presence of \ceFe2+/\ceFe3+\ce{Fe^{2+}} / \ce{Fe^{3+}} pairs enables electrical conduction through small hopping, where charge carriers (holes) migrate between iron sites, rendering a p-type with enhanced conductivity compared to the hypothetical stoichiometric form. Additionally, the cation vacancies reduce the overall below that of an ideal \ceFeO\ce{FeO} crystal, with measured values around 5.75 g/cm³ for typical non-stoichiometric compositions. In the of the Fe-O system, is thermodynamically stable only above approximately 575 °C, where the non-stoichiometric phase field exists between the iron metal and higher oxides; below this temperature, it decomposes into iron and \ceFe3O4\ce{Fe_3O_4}, though quenched samples remain metastable at . The was named after German metallurgist Wüst (1860–1938), who first synthesized it, recognizing its distinct non-stoichiometric character in natural and artificial systems.

Preparation

Laboratory synthesis

Iron(II) oxide can be synthesized in the laboratory through the thermal decomposition of iron(II) oxalate dihydrate under an inert atmosphere to prevent oxidation. The reaction, represented as \ceFeC2O4⋅2H2O−>[600−800∘C]FeO+CO+CO2+2H2O\ce{FeC2O4 \cdot 2H2O ->[600-800^\circ C] FeO + CO + CO2 + 2H2O}, yields black FeO upon rapid cooling to avoid disproportionation into metallic iron and magnetite. This method produces relatively pure samples when conducted in a stream of CO₂ or nitrogen at temperatures around 600–800 °C in a quartz vessel. Another approach involves the controlled oxidation of metallic iron by heating fine iron powder or wire in a limited oxygen environment at elevated temperatures, typically 400–600 °C, to form a surface layer of FeO before further oxidation to higher oxides occurs. This technique requires precise control of oxygen to halt the process at the stage, often achieved in a with a controlled gas mixture. A common method is the controlled reduction of higher iron oxides, such as (Fe₂O₃) or (Fe₃O₄), using (H₂) or (CO) at temperatures of 400–600 °C under an inert atmosphere. The reaction proceeds stepwise: \ceFe2O3+3H2−>2FeO+3H2O\ce{Fe2O3 + 3H2 -> 2FeO + 3H2O} or \ceFe2O3+3CO−>2FeO+3CO2\ce{Fe2O3 + 3CO -> 2FeO + 3CO2}, with conditions adjusted (e.g., low reducing gas partial pressure) to stop at the FeO stage and prevent further reduction to metallic iron. To obtain stoichiometric FeO, a high-pressure method entails heating non-stoichiometric Fe_{0.95}O with excess metallic iron at 770 °C and pressures exceeding 36 kbar, resulting in the diffusion of iron into the lattice to achieve the exact Fe:O ratio of 1:1. This synthesis, first demonstrated in piston-cylinder apparatus, addresses the inherent cation vacancies in ambient-pressure FeO. Purification of laboratory-synthesized FeO often involves hydrogen reduction to eliminate traces of higher iron oxides like Fe_{3}O_{4}, where the sample is heated in a hydrogen stream at 500–700 °C, selectively reducing magnetite back to wüstite without further reduction to iron. Vacuum distillation under reduced pressure can also remove volatile impurities or adsorbed species, though it is less common due to FeO's low volatility and is typically applied post-synthesis heating.

Industrial production

Iron(II) oxide, commonly referred to as , is primarily produced on an industrial scale as an intermediate byproduct during the reduction of in blast furnaces for . In this process, , mainly (Fe₂O₃), undergoes stepwise reduction by (CO) and (H₂) gases generated from coke . The sequence progresses from to (Fe₃O₄) in the upper furnace, then to (FeO) in the middle to lower zones at temperatures of approximately 900–1200°C, before final reduction to metallic iron. The formation of occurs via reactions such as Fe₃O₄ + CO → 3FeO + CO₂, with the material exhibiting non-stoichiometry (Fe_{1-x}O, where 0.05 < x < 0.16) due to cation vacancies that enhance its reducibility. An alternative method involves the of iron sponge () or fine iron powder in controlled atmospheres to selectively form without progressing to higher oxides like or . This is achieved at temperatures above 570°C using limited oxidizing agents such as or CO₂, where the oxygen is maintained low to stabilize the FeO phase in equilibrium. For instance, in demonstration-scale reactors for or production, oxidation degrees up to 90% can be reached with CO₂ flow rates 5–6 times the stoichiometric requirement, utilizing commercial iron pellets containing residual phases. In all these processes, the resulting iron(II) oxide is typically non-stoichiometric with variable oxygen content (Fe_{0.84}O to Fe_{0.95}O) and contains impurities like silica or alumina from raw materials, achieving purities around 85–95% FeO. It is often used directly without purification due to economic considerations, with yields in intermediates approaching theoretical limits based on input, while salt-based methods yield 70–80% based on salt conversion. Higher purity variants may reference scaled techniques but are not standard in high-volume production.

Reactions

Thermal stability and decomposition

Iron(II) oxide, also known as wüstite (FeO), is thermodynamically unstable below approximately 570 °C and undergoes via the reaction 4FeO→Fe+Fe3O44 \mathrm{FeO} \to \mathrm{Fe} + \mathrm{Fe_3O_4}. This process occurs eutectoidally, forming a of iron and , but the phase remains quenchable and metastable at due to slow kinetics that inhibit spontaneous decomposition under ambient conditions. The of solid FeO is ΔHf∘=−272 kJ/mol\Delta H_f^\circ = -272 \, \mathrm{kJ/mol}, indicating moderate exothermicity in its synthesis from elements but highlighting its propensity for relative to more stable iron oxides like Fe₃O₄. Decomposition kinetics of wüstite are diffusion-controlled, with an of about 33 kJ/mol during heating up to 823 K, beyond which regeneration to the wüstite phase becomes thermodynamically favored. Under high pressure, wüstite transitions from its cubic rock-salt (B1) structure to a hexagonal NiAs-type (B8) structure at approximately 100 GPa and 300 K, marking a shift toward metallic behavior.

Redox reactions

Iron(II) oxide undergoes oxidation in the presence of oxygen to form magnetite, according to the reaction 6FeO+O2→2Fe3O46 \mathrm{FeO} + \mathrm{O_2} \rightarrow 2 \mathrm{Fe_3O_4}. This process involves the partial oxidation of Fe(II) to a mixed Fe(II)/Fe(III) state and is thermodynamically favorable at elevated temperatures above 570 °C, where wüstite (FeO) is stable. Reduction of iron(II) oxide can be achieved using gas, following the FeO+H2→Fe+H2O\mathrm{FeO} + \mathrm{H_2} \rightarrow \mathrm{Fe} + \mathrm{H_2O}. This reaction is exothermic and commonly employed in reduction processes for iron production, with kinetics enhanced at temperatures around 800–1000 °C under atmospheres. In metallurgical applications, serves as an alternative reductant via FeO+CO→Fe+CO2\mathrm{FeO} + \mathrm{CO} \rightarrow \mathrm{Fe} + \mathrm{CO_2}, facilitating iron extraction in blast furnaces. Iron(II) oxide reacts with halogens such as chlorine to produce iron(II) chloride and oxygen, as in 2FeO+Cl2→2FeCl2+O22 \mathrm{FeO} + \mathrm{Cl_2} \rightarrow 2 \mathrm{FeCl_2} + \mathrm{O_2}. This redox reaction proceeds at high temperatures and demonstrates the oxidizing power of halogens toward the oxide, liberating oxygen while reducing Cl₂ to FeCl₂. In electrochemical contexts, the reduction of iron(II) from FeO relates to the standard reduction potential of the Fe²⁺/Fe couple, which is -0.44 V versus the standard hydrogen electrode. This value indicates that FeO reduction to metallic iron requires a sufficiently negative potential, influencing its behavior in electrolytic processes or corrosion scenarios involving dissolved Fe²⁺ species derived from the oxide. The kinetics of reactions involving iron(II) oxide are influenced by and atmospheric conditions. Smaller particles exhibit higher initial oxidation rates due to increased surface area, with exponential enhancement observed for sizes below 40 nm during oxygen exposure. Reduction rates with similarly accelerate in finer particles and under pure H₂ atmospheres, as opposed to mixed gases, highlighting the role of and surface reactivity.

Occurrence

In Earth's interior

Iron(II) oxide, primarily in the form of ferropericlase ((Mg,Fe)O), constitutes approximately 20% of the volume of Earth's , making it the second most abundant mineral phase after bridgmanite. This incorporates iron in the divalent Fe²⁺ state, with typical iron content ranging from 10 to 25 mol%, depending on local conditions and bulk composition. In pyrolitic models of mantle composition, ferropericlase forms a significant oxide component alongside silicate phases, influencing the overall and of the region from about 660 km to 2900 km depth. Ferropericlase forms in the under extreme high-pressure and reducing conditions through the equilibrium transformation of silicates, such as the dehydration or of hydrous phases like during . These conditions, characterized by pressures exceeding 23 GPa and temperatures around 1500–2500°C, favor the partitioning of iron and magnesium into the oxide phase separate from silicon-bearing minerals. The reducing environment, maintained by low oxygen from the mantle's bulk composition, ensures the dominance of Fe²⁺ over higher oxidation states, preventing incorporation into ferric forms that would alter phase stability. The stability of ferropericlase in the is extended by gigapascal-level pressures, which suppress the decomposition or phase transitions observed at lower pressures; for instance, pure FeO () is only stable above approximately 570°C at , but mantle conditions stabilize the rock-salt structure to much broader temperature ranges. This persistence contributes to the mineral's role in deep mantle dynamics, as its B1 rock-salt structure (with minor distortions under pressure) accommodates variable iron content without disrupting the overall assemblage. Geophysically, ferropericlase influences lower mantle models through its contributions to electrical conductivity and seismic wave propagation. The mineral's iron content enhances electrical conductivity, potentially by up to several orders of magnitude compared to pure MgO, which helps explain observed anomalies in geomagnetic field variations and supports models of core-mantle boundary interactions. Additionally, the spin crossover of Fe²⁺ from high- to low-spin states at mid-mantle depths (around 1000–2000 km) induces subtle changes in density and elasticity, producing detectable heterogeneities in seismic velocities that refine tomographic images of mantle convection and subducted slabs.

In meteorites and other extraterrestrial materials

Iron(II) oxide, commonly known as wüstite (FeO), occurs in various s as a minor phase resulting from oxidation processes during formation or alteration. In iron meteorites of the Aletai , such as the Akebulake meteorite, wüstite is identified as a associated with native iron and other oxides, forming under reducing conditions close to the iron-wüstite (IW) oxygen buffer. In chondritic meteorites, particularly CR chondrites, wüstite appears in the matrix as a scarce phase alongside ferrolizardite and sulfides, indicating localized oxidation in the parent body. These occurrences highlight wüstite's role as an oxidation product in , often metastable and preserved due to rapid cooling or low-temperature environments. Beyond meteorites, wüstite is implicated in the interiors of planetary bodies like Mars and the Moon based on spectroscopic and geochemical analyses. For the Moon, mantle oxygen fugacity estimates from volcanic glasses and basalts place conditions approximately 1 log unit below the IW buffer, within the stability field of wüstite, suggesting its potential presence in deep lunar silicates. On Mars, core-mantle differentiation models indicate oxygen fugacities slightly more oxidizing than the IW buffer, consistent with wüstite formation during early accretion, though direct detection relies on inferred redox states from meteoritic analogs and orbital spectroscopy showing iron oxide signatures in surface-derived materials. In stellar environments, iron(II) oxide contributes to dust formation in supernova ejecta and circumstellar envelopes. Observations of a 21 μm emission feature in spectra from evolved stars and supernova remnants suggest the presence of FeO nanoparticles, likely condensing from cooling gas in oxygen-rich environments. In planetary nebulae, similar oxide grains may form via gas-phase condensation, enriching dust with iron oxides. Recent studies in the 2020s have explored wüstite's properties in models, particularly its electrical conductivity under high-pressure, high-temperature conditions. Experiments on molten iron oxides reveal redox-dependent structural variations, with iron dominating in FeO5 and FeO4 units under reducing conditions, informing models of interior dynamics in rocky s where wüstite-like phases influence and magnetic field generation.

Uses

As a pigment

Iron(II) oxide, known for its black crystalline form, serves as a stable in various artistic and industrial applications, particularly in ceramics where it produces durable black hues during reduction firings. It is also employed in glass coloring to achieve deep black tones and in paints for its permanence and resistance to fading. These properties stem from its strong fluxing action in high-temperature processes, enhancing color stability without compromising material integrity. Historically, black iron oxide, including forms, has been utilized since antiquity in glazes to create enduring black decorations on ceramics, as evidenced by ancient artifacts from various civilizations. This long-standing application underscores its reliability as a in traditional artisanal techniques. Synthetic iron oxides are FDA-approved as color additives for use in , including those applied to the eye area, due to their safety profile and exemption from certification requirements; iron(II) oxide is used in some cosmetic formulations. It finds application in tattoo inks for black pigmentation, offering color stability and in permanent cosmetics. Iron oxides are also permitted in pharmaceuticals as pigments in certain formulations. To achieve varied tones, iron(II) oxide is often blended with other iron oxides, such as (Fe₂O₃) or (FeOOH), allowing for customizable shades from pure to nuanced browns in pigment formulations. This mixing enables precise control over hue in applications like paints and glazes.

In environmental and other applications

In , iron(II) oxide acts as a precursor for iron-based catalysts in Fischer-Tropsch synthesis, where it is reduced to metallic iron under atmospheres, facilitating the conversion of (CO and H₂) to hydrocarbons with selectivities favoring longer-chain products when promoted with alkali metals. The reduction pathway typically proceeds from higher oxides through FeO to active metallic phases, enabling industrial-scale production of fuels and chemicals. Iron(II) oxide shows promise in as a component in iron-air batteries, where its reversible between Fe and FeO contributes to the reaction, supporting long-duration discharge cycles with theoretical energy densities around 1,000 Wh/kg and cycle efficiencies exceeding 80% in solid-state configurations. This positions it as a low-cost, sustainable option for grid-scale storage due to abundant raw materials and minimal environmental impact. Handling precautions for iron(II) oxide include avoiding ignition sources, as the material is flammable and may ignite spontaneously in air, particularly in finely divided forms, with combustibility risks increasing above 200 °C. The classifies it without specific hazards under REACH, but recommends dust control and to prevent or skin contact.

References

  1. https://pubchem.ncbi.nlm.nih.gov/compound/Ferrous-oxide
  2. https://pubchem.ncbi.nlm.nih.gov/compound/Wustite
  3. https://www.americanelements.com/iron-ii-oxide-1345-25-1
  4. https://webbook.nist.gov/cgi/inchi?ID=C1345251&Type=JANAFS&Table=on
  5. https://arxiv.org/abs/1203.1201
  6. https://www.chemicalbook.com/ChemicalProductProperty_EN_CB6715199.htm
  7. https://reade.com/product/iron-ii-oxide-feo/
  8. https://pubs.aip.org/aip/apl/article/108/4/041606/31528/Growth-and-magnetic-properties-of-ultrathin
  9. https://pubs.usgs.gov/wsp/1459a/report.pdf
  10. https://www.researchgate.net/figure/Solubility-of-FeII-vs-pH-as-controlled-by-the-solubility-of-FeOH-2-s-a-The_fig2_355110583
  11. https://next-gen.materialsproject.org/materials/mp-18905
  12. https://rruff.geo.arizona.edu/AMS/minerals/Wustite
  13. https://www.mindat.org/min-4316.html
  14. https://pubs.aip.org/aip/jcp/article/153/19/194702/199277/Covalency-in-Fe2O3-and-FeO-Consequences-for-XPS
  15. https://www.nature.com/articles/s43246-020-00080-4
  16. https://agupubs.onlinelibrary.wiley.com/doi/abs/10.1029/RG022i001p00037
  17. https://www.sciencedirect.com/science/article/abs/pii/S0079678618300281
  18. https://pubs.geoscienceworld.org/minersoc/books/edited-volume/942/chapter/106818628/Non-stoichiometry-defects-and-superstructures-in
  19. https://www.nature.com/articles/s41524-025-01527-3
  20. https://www.sciencedirect.com/science/article/abs/pii/0022369786900053
  21. https://www.pearson.com/channels/general-chemistry/textbook-solutions/mcmurry-8th-edition-9781292336145/ch-12-solids-and-solid-state-materials/the-mineral-wustite-is-a-nonstoichiometric-iron-oxide-with-the-empirical-formula-1
  22. https://www.sciencedirect.com/science/article/abs/pii/S0169433225007457
  23. https://www.sciencedirect.com/science/article/abs/pii/0022190267802021
  24. https://pubs.rsc.org/en/content/articlelanding/2006/jm/b514565a
  25. https://chemistry.stackexchange.com/questions/118606/what-is-the-most-convenient-way-to-prepare-ferrous-oxide-feo-in-the-laboratory
  26. https://onlinelibrary.wiley.com/doi/10.1002/srin.201900108
  27. https://pubs.aip.org/aip/jcp/article/47/11/4559/85273/High-Pressure-Synthesis-of-the-Stoichiometric
  28. https://hal.science/hal-00265636v1/document
  29. https://www.ispatguru.com/chemistry-of-the-ironmaking-by-blast-furnace-process/
  30. https://www.intechopen.com/chapters/1126784
  31. https://onlinelibrary.wiley.com/doi/full/10.1002/srin.202400891
  32. https://deepblue.lib.umich.edu/bitstream/handle/2027.42/30621/0000262.pdf?sequence=1
  33. https://www.tandfonline.com/doi/full/10.1080/21663831.2019.1684395
  34. https://webbook.nist.gov/cgi/inchi?ID=C1345251&Mask=2
  35. https://digitalcommons.mtu.edu/michigantech-p/3398/
  36. https://www.researchgate.net/figure/Phase-diagram-of-wuestite-with-comparison-to-electrical-resistivity-measurements-Melting_fig1_228444269
  37. https://www.chemicalaid.com/tools/equationbalancer.php?equation=FeO%2B%252B%2BO2%2B%253D%2BFe3O4&hl=en
  38. https://www.chemicalaid.com/tools/equationbalancer.php?equation=FeO%2B%252B%2BH2%2B%253D%2BFe%2B%252B%2BH2O&hl=en
  39. https://www.chemicalaid.com/tools/equationbalancer.php?equation=FeO%2B%252B%2BCl2%2B%253D%2BFeCl2%2B%252B%2BO2&hl=en
  40. https://chem.libretexts.org/Ancillary_Materials/Reference/Reference_Tables/Electrochemistry_Tables/P2%253A_Standard_Reduction_Potentials_by_Value
  41. https://www.nature.com/articles/srep18818
  42. https://www.sciencedirect.com/science/article/pii/S000925091400428X
  43. https://link.aps.org/doi/10.1103/PhysRevLett.114.117202
  44. https://www.pnas.org/doi/10.1073/pnas.0508919102
  45. https://agupubs.onlinelibrary.wiley.com/doi/full/10.1002/2016JB013352
  46. https://www.science.org/doi/10.1126/sciadv.adt9106
  47. https://www.nature.com/articles/s41467-022-35110-x
  48. https://www.sciencedirect.com/science/article/abs/pii/S0012821X04002018
  49. https://www.pnas.org/doi/10.1073/pnas.252782399
  50. https://agupubs.onlinelibrary.wiley.com/doi/full/10.1029/2007gl030523
  51. https://www.pnas.org/doi/abs/10.1073/pnas.1322427111
  52. https://www.hou.usra.edu/meetings/lpsc2015/pdf/2561.pdf
  53. https://pubs.geoscienceworld.org/msa/rimg/article-abstract/68/1/493/87595/Redox-Conditions-on-Small-Bodies-the-Moon-and-Mars
  54. https://agupubs.onlinelibrary.wiley.com/doi/full/10.1002/jgre.20093
  55. https://iopscience.iop.org/article/10.3847/1538-4357/aa7deb
  56. https://www.frontiersin.org/journals/astronomy-and-space-sciences/articles/10.3389/fspas.2022.908217/full
  57. https://doi.org/10.1038/s43246-020-00080-4
  58. https://www.theceramicshop.com/product/365/iron-oxide-black/
  59. https://cameo.mfa.org/wiki/Ferrous_oxide
  60. https://digitalfire.com/oxide/feo
  61. https://hfpappexternal.fda.gov/scripts/fdcc/index.cfm?set=ColorAdditives&id=IronOxideSynthetic
  62. https://www.drugs.com/inactive/ferrous-oxide-389.html
  63. https://www.golchhaoxides.com/iron-oxide-pigments.php
  64. https://www.wetcanvas.com/forums/topic/iron-oxide-pigments/
  65. https://pubs.acs.org/doi/10.1021/jp010288u
  66. https://www.sciencedirect.com/science/article/abs/pii/S2095495620301583
  67. https://www.sciencedirect.com/science/article/abs/pii/S0306261918315757
  68. https://pubs.rsc.org/en/content/articlehtml/2024/se/d4se01224k
  69. https://nj.gov/health/eoh/rtkweb/documents/fs/1036.pdf
  70. https://echa.europa.eu/substance-information/-/substanceinfo/100.014.292
Add your contribution
Related Hubs
User Avatar
No comments yet.